# data[].attributes.body
Form by crystallzation
Two types: Intrusive and Extrusive
Textures include vesicular,glassy,aphanitic, and phaneritic
Compositions include:mafic or balsaltic rocks and felsic or granite rocks
Igneous rock features: batholith, dike, sill, caldera

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Non metallic
Has copper coloring
Produces white streak
Cleavage
Can scratch a penny
Find in ceramic glass

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Metallic
Has dark grey coloring
Produces dark grey streak
Fracture
Can scratch a penny
Found in ore of iron

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Unit cell of magnetite. The gray spheres are oxygen, green are divalent iron, blue are trivalent iron. Also shown are an iron atom in an octahedral space (light blue) and another in a tetrahedral space (gray).

Magnetite is a mineral and one of the main iron ores, with the chemical formula . It is one of the oxides of iron, and is ferrimagnetic; it is attracted to a magnet and can be magnetized to become a permanent magnet itself. With the exception of extremely rare native iron deposits, it is the most magnetic of all the naturally occurring minerals on Earth. Naturally magnetized pieces of magnetite, called lodestone, will attract small pieces of iron, which is how ancient peoples first discovered the property of magnetism.

Magnetite is black or brownish-black with a metallic luster, has a Mohs hardness of 5-6 and leaves a black streak. Small grains of magnetite are very common in igneous and metamorphic rocks.

The chemical IUPAC name is iron(II,III) oxide and the common chemical name is ferrous-ferric oxide.

In addition to igneous rocks, magnetite also occurs in sedimentary rocks, including banded iron formations and in lake and marine sediments as both detrital grains and as magnetofossils. Magnetite nanoparticles are also thought to form in soils, where they probably oxidize rapidly to maghemite.

The chemical composition of magnetite is Fe ²⁺(Fe ³⁺) ₂(O ²⁻) ₄. This indicates that magnetite contains both ferrous ( divalent) and ferric ( trivalent) iron, suggesting crystallization in an environment containing intermediate levels of oxygen. The main details of its structure were established in 1915. It was one of the first crystal structures to be obtained using X-ray diffraction. The structure is inverse spinel, with O 2− ions forming a face-centered cubic lattice and iron cations occupying interstitial sites. Half of the Fe ³⁺ cations occupy tetrahedral sites while the other half, along with Fe ²⁺ cations, occupy octahedral sites. The unit cell consists of 32O 2− ions and unit cell length is a = 0.839 nm.

As a member of the inverse spinel group, magnetite can form solid solutions with similarly structured minerals, including ulvospinel () and magnesioferrite ().

Titanomagnetite, also known as titaniferous magnetite, is a solid solution between magnetite and ulvospinel that crystallizes in many mafic igneous rocks. Titanomagnetite may undergo oxy-exsolution during cooling, resulting in ingrowths of magnetite and ilmenite.

Natural and synthetic magnetite occurs most commonly as octahedral crystals bounded by {111} planes and as rhombic-dodecahedra. Twinning occurs on the {111} plane.

Hydrothermal synthesis usually produces single octahedral crystals which can be as large as 10 mm (0.39 in) across. In the presence of mineralizers such as 0.1M HI or 2M and at 0.207 MPa at 416-800 °C, magnetite grew as crystals whose shapes were a combination of rhombic-dodechahedra forms. The crystals were more rounded than usual. The appearance of higher forms was considered as a result from a decrease in the surface energies caused by the lower surface to volume ratio in the rounded crystals.

Magnetite has been important in understanding the conditions under which rocks form. Magnetite reacts with oxygen to produce hematite, and the mineral pair forms a buffer that can control how oxidizing its environment is (the oxygen fugacity). This buffer is known as the hematite-magnetite or HM buffer. At lower oxygen levels, magnetite can form a buffer with quartz and fayalite known as the QFM buffer. At still lower oxygen levels, magnetite forms a buffer with wüstite known as the MW buffer. The QFM and MW buffers have been used extensively in laboratory experiments on rock chemistry. The QFM buffer, in particular, produces an oxygen fugacity close to that of most igneous rocks.

Commonly, igneous rocks contain solid solutions of both titanomagnetite and hemoilmenite or titanohematite. Compositions of the mineral pairs are used to calculate oxygen fugacity: a range of oxidizing conditions are found in magmas and the oxidation state helps to determine how the magmas might evolve by fractional crystallization. Magnetite also is produced from peridotites and dunites by serpentinization.

Lodestones were used as an early form of magnetic compass. Magnetite has been a critical tool in paleomagnetism, a science important in understanding plate tectonics and as historic data for magnetohydrodynamics and other scientific fields.

The relationships between magnetite and other iron oxide minerals such as ilmenite, hematite, and ulvospinel have been much studied; the reactions between these minerals and oxygen influence how and when magnetite preserves a record of the Earth's magnetic field.

At low temperatures, magnetite undergoes a crystal structure phase transition from a monoclinic structure to a cubic structure known as the Verwey transition. Optical studies show that this metal to insulator transition is sharp and occurs around 120K. The Verwey transition is dependent on grain size, domain state, pressure, and the iron-oxygen stoichiometry. An isotropic point also occurs near the Verwey transition around 130K, at which point the sign of the magnetocrystalline anisotropy constant changes from positive to negative. The Curie temperature of magnetite is 580 °C (853 K; 1,076 °F).

If magnetite is in a large enough quantity it can be found in aeromagnetic surveys using a magnetometer which measures magnetic intensities.

Solid magnetite particles melt at about 1,583-1,597 °C (2,881-2,907 °F).

Magnetite and other heavy minerals (dark) in a quartz beach sand (Chennai, India).

Magnetite is sometimes found in large quantities in beach sand. Such black sands (mineral sands or iron sands) are found in various places, such as Lung Kwu Tan of Hong Kong; California, United States; and the west coast of the North Island of New Zealand. The magnetite, eroded from rocks, is carried to the beach by rivers and concentrated by wave action and currents. Huge deposits have been found in banded iron formations. These sedimentary rocks have been used to infer changes in the oxygen content of the atmosphere of the Earth.

Large deposits of magnetite are also found in the Atacama region of Chile ( Chilean Iron Belt); the Valentines region of Uruguay; Kiruna, Sweden; the Tallawang Region of New South Wales; and in the Adirondack region of New York in the United States. Kediet ej Jill, the highest mountain of Mauritania, is made entirely of the mineral. Deposits are also found in Norway, Romania, and Ukraine. Magnetite-rich sand dunes are found in southern Peru. In 2005, an exploration company, Cardero Resources, discovered a vast deposit of magnetite-bearing sand dunes in Peru. The dune field covers 250 square kilometers (100 sq mi), with the highest dune at over 2,000 meters (6,560 ft) above the desert floor. The sand contains 10% magnetite.

In large enough quantities magnetite can affect compass navigation. In Tasmania there are many areas with highly magnetized rocks that can greatly influence compasses. Extra steps and repeated observations are required when using a compass in Tasmania to keep navigation problems to the minimum.

Magnetite crystals with a cubic habit are rare but have been found at Balmat, St. Lawrence County, New York, and at Långban, Sweden. This habit may be a result of crystallization in the presence of cations such as zinc.

Magnetite can also be found in fossils due to biomineralization and are referred to as magnetofossils. There are also instances of magnetite with origins in space coming from meteorites.

Biomagnetism is usually related to the presence of biogenic crystals of magnetite, which occur widely in organisms. These organisms range from magnetotactic bacteria (e.g., Magnetospirillum magnetotacticum) to animals, including humans, where magnetite crystals (and other magnetically sensitive compounds) are found in different organs, depending on the species. Biomagnetites account for the effects of weak magnetic fields on biological systems. There is also a chemical basis for cellular sensitivity to electric and magnetic fields ( galvanotaxis).

Magnetite magnetosomes in Gammaproteobacteria

Pure magnetite particles are biomineralized in magnetosomes, which are produced by several species of magnetotactic bacteria. Magnetosomes consist of long chains of oriented magnetite particle that are used by bacteria for navigation. After the death of these bacteria, the magnetite particles in magnetosomes may be preserved in sediments as magnetofossils. Some types of anaerobic bacteria that are not magnetotactic can also create magnetite in oxygen free sediments by reducing amorphic ferric oxide to magnetite.

Several species of birds are known to incorporate magnetite crystals in the upper beak for magnetoreception, which (in conjunction with cryptochromes in the retina) gives them the ability to sense the direction, polarity, and magnitude of the ambient magnetic field.

Chitons, a type of mollusk, have a tongue-like structure known as a radula, covered with magnetite-coated teeth, or denticles. The hardness of the magnetite helps in breaking down food.

Biological magnetite may store information about the magnetic fields the organism was exposed to, potentially allowing scientists to learn about the migration of the organism or about changes in the Earth's magnetic field over time.

Living organisms can produce magnetite. In humans, magnetite can be found in various parts of the brain including the frontal, parietal, occipital, and temporal lobes, brainstem, cerebellum and basal ganglia. Iron can be found in three forms in the brain - magnetite, hemoglobin (blood) and ferritin (protein), and areas of the brain related to motor function generally contain more iron. Magnetite can be found in the hippocampus. The hippocampus is associated with information processing, specifically learning and memory. However, magnetite can have toxic effects due to its charge or magnetic nature and its involvement in oxidative stress or the production of free radicals. Research suggests that beta-amyloid plaques and tau proteins associated with neurodegenerative disease frequently occur after oxidative stress and the build-up of iron.

Some researchers also suggest that humans possess a magnetic sense, proposing that this could allow certain people to use magnetoreception for navigation. The role of magnetite in the brain is still not well understood, and there has been a general lag in applying more modern, interdisciplinary techniques to the study of biomagnetism.

Electron microscope scans of human brain-tissue samples are able to differentiate between magnetite produced by the body's own cells and magnetite absorbed from airborne pollution, the natural forms being jagged and crystalline, while magnetite pollution occurs as rounded nanoparticles. Potentially a human health hazard, airborne magnetite is a result of pollution (specifically combustion). These nanoparticles can travel to the brain via the olfactory nerve, increasing the concentration of magnetite in the brain. In some brain samples, the nanoparticle pollution outnumbers the natural particles by as much as 100:1, and such pollution-borne magnetite particles may be linked to abnormal neural deterioration. In one study, the characteristic nanoparticles were found in the brains of 37 people: 29 of these, aged 3 to 85, had lived and died in Mexico City, a significant air pollution hotspot. Some of the further eight, aged 62 to 92, from Manchester, England, had died with varying severities of neurodegenerative diseases. Such particles could conceivably contribute to diseases like Alzheimer's disease. Though a causal link has not yet been established, laboratory studies suggest that iron oxides such as magnetite are a component of protein plaques in the brain. Such plaques have been linked to Alzheimer's disease.

Increased iron levels, specifically magnetic iron, have been found in portions of the brain in Alzheimer's patients. Monitoring changes in iron concentrations may make it possible to detect the loss of neurons and the development of neurodegenerative diseases prior to the onset of symptoms due to the relationship between magnetite and ferritin. In tissue, magnetite and ferritin can produce small magnetic fields which will interact with magnetic resonance imaging (MRI) creating contrast. Huntington patients have not shown increased magnetite levels; however, high levels have been found in study mice.

Due to its high iron content, magnetite has long been a major iron ore. It is reduced in blast furnaces to pig iron or sponge iron for conversion to steel.

Audio recording using magnetic acetate tape was developed in the 1930s. The German magnetophon first utilized magnetite powder that BASF coated onto cellulose acetate before soon switching to gamma ferric oxide for its superior morphology. Following World War II, 3M Company continued work on the German design. In 1946, the 3M researchers found they could also improve their own magnetite-based paper tape, which utilized powders of cubic crystals, by replacing the magnetite with needle-shaped particles of gamma ferric oxide (γ-Fe ₂O ₃).

Approximately 2-3% of the world's energy budget is allocated to the Haber Process for nitrogen fixation, which relies on magnetite-derived catalysts. The industrial catalyst is obtained from finely ground iron powder, which is usually obtained by reduction of high-purity magnetite. The pulverized iron metal is burnt (oxidized) to give magnetite or wüstite of a defined particle size. The magnetite (or wüstite) particles are then partially reduced, removing some of the oxygen in the process. The resulting catalyst particles consist of a core of magnetite, encased in a shell of wüstite, which in turn is surrounded by an outer shell of iron metal. The catalyst maintains most of its bulk volume during the reduction, resulting in a highly porous high-surface-area material, which enhances its effectiveness as a catalyst.

Magnetite micro- and nanoparticles are used in a variety of applications, from biomedical to environmental. One use is in water purification: in high gradient magnetic separation, magnetite nanoparticles introduced into contaminated water will bind to the suspended particles (solids, bacteria, or plankton, for example) and settle to the bottom of the fluid, allowing the contaminants to be removed and the magnetite particles to be recycled and reused. This method works with radioactive and carcinogenic particles as well, making it an important cleanup tool in the case of heavy metals introduced into water systems.

Another application of magnetic nanoparticles is in the creation of ferrofluids. These are used in several ways. Ferrofluids can be used for targeted drug delivery in the human body. The magnetization of the particles bound with drug molecules allows "magnetic dragging" of the solution to the desired area of the body. This would allow the treatment of only a small area of the body, rather than the body as a whole, and could be highly useful in cancer treatment, among other things. Ferrofluids are also used in magnetic resonance imaging (MRI) technology.

For the separation of coal from waste, dense medium baths were used. This technique employed the difference in densities between coal (1.3-1.4 tonnes per m ³) and shales (2.2-2.4 tonnes per m ³). In a medium with intermediate density (water with magnetite), stones sank and coal floated.

Magnetene is a 2 dimensional flat sheet of magnetite noted for its ultra-low-friction behavior.

1. ^ Anthony, John W.; Bideaux, Richard A.; Bladh, Kenneth W. "Magnetite" (PDF). Handbook of Mineralogy. Chantilly, VA: Mineralogical Society of America. p. 333.

2. "Magnetite". mindat.org and the Hudson Institute of Mineralogy.

3. ^ Jacobsen, S.D.; Reichmann, H.J.; Kantor, A.; Spetzler, H.A. (2005). "A gigahertz ultrasonic interferometer for the diamond anvil cell and high-pressure elasticity of some iron-oxide minerals". In Chen, J.; Duffy, T.S.; Dobrzhinetskaya, L.F.; Wang, Y.; Shen, G. (eds.). Advances in High-Pressure Technology for Geophysical Applications. Elsevier Science. pp. 25-48. doi:10.1016/B978-044451979-5.50004-1. ISBN 978-0-444-51979-5.

4. ^ Hurlbut, Cornelius Searle; W. Edwin Sharp; Edward Salisbury Dana (1998). Dana's minerals and how to study them. John Wiley and Sons. pp. 96. ISBN 978-0-471-15677-2.

5. ^ Harrison, R. J.; Dunin-Borkowski, RE; Putnis, A (2002). "Direct imaging of nanoscale magnetic interactions in minerals". Proceedings of the National Academy of Sciences. 99 (26): 16556-16561. Bibcode:2002PNAS...9916556H. doi:10.1073/pnas.262514499. PMC . PMID 12482930.

6. ^ Du Trémolet de Lacheisserie, Étienne; Damien Gignoux; Michel Schlenker (2005). Magnetism: Fundamentals. Springer. pp. 3-6. ISBN 0-387-22967-1.

7. ^ Kesler, Stephen E.; Simon, Adam F. (2015). Mineral resources, economics and the environment (2nd ed.). Cambridge, United Kingdom: Cambridge University Press. ISBN 9781107074910. OCLC 907621860.

8. ^ an alternative visualisation of the crystal structure of Magnetite using JSMol is found here.

9. ^ Gasparov, L. V.; et al. (2000). "Infrared and Raman studies of the Verwey transition in magnetite". Physical Review B. 62 (12): 7939. arXiv:cond-mat/9905278. Bibcode:2000PhRvB..62.7939G. CiteSeerX 10.1.1.242.6889. doi:10.1103/PhysRevB.62.7939. S2CID 39065289.

10. ^ Gasparov, L. V.; et al. (2005). "Magnetite: Raman study of the high-pressure and low-temperature effects". Journal of Applied Physics. 97 (10): 10A922. arXiv:. Bibcode:2005JAP....97jA922G. doi:10.1063/1.1854476. S2CID 55568498. 10A922.

11. ^ Gubbins, D.; Herrero-Bervera, E., eds. (2007). Encyclopedia of geomagnetism and paleomagnetism. Springer Science & Business Media.

12. ^ Fabian, K.; Shcherbakov, V. P.; McEnroe, S. A. (April 2013). "Measuring the Curie temperature". Geochemistry, Geophysics, Geosystems. 14 (4): 947-961. Bibcode:2013GGG....14..947F. doi:10.1029/2012GC004440.

13. ^ Wallace, Roberts M. (1976). "Geological reconnaissance of some Uruguayan iron and manganese deposits in 1962" (PDF). U.S. Geological Survey Open File Report. Open-File Report. 76-466. doi:10.3133/ofr76466.

14. ^ Knipping, Jaayke L.; Bilenker, Laura D.; Simon, Adam C.; Reich, Martin; Barra, Fernando; Deditius, Artur P.; Lundstrom, Craig; Bindeman, Ilya; Munizaga, Rodrigo (July 2015). "Giant Kiruna-type deposits form by efficient flotation of magmatic magnetite suspensions". Geology. 43 (7): 591-594. Bibcode:2015Geo....43..591K. doi:10.1130/G36650.1. hdl:10533/228146.

15. ^ Clark, David A. (September 2012). "Interpretation of the magnetic gradient tensor and normalized source strength applied to the Tallawang magnetite skarn deposit, New South Wales, Australia". SEG Technical Program Expanded Abstracts 2012: 1-5. doi:10.1190/segam2012-0700.1.

16. ^ Valley, Peter M.; Hanchar, John M.; Whitehouse, Martin J. (April 2011). "New insights on the evolution of the Lyon Mountain Granite and associated Kiruna-type magnetite-apatite deposits, Adirondack Mountains, New York State". Geosphere. 7 (2): 357-389. Bibcode:2011Geosp...7..357V. doi:10.1130/GES00624.1.

17. ^ Leaman, David. "Magnetic Rocks - Their Effect on Compass Use and Navigation in Tasmania" (PDF). Archived from the original (PDF) on 2017-03-29.

18. ^ Boström, Kurt (15 December 1972). "Magnetite Crystals of Cubic Habit from Långban, Sweden". Geologiska Föreningen i Stockholm Förhandlingar. 94 (4): 572-574. doi:10.1080/11035897209453690.

19. ^ Chang, S. B. R.; Kirschvink, J. L. (May 1989). "Magnetofossils, the Magnetization of Sediments, and the Evolution of Magnetite Biomineralization" (PDF). Annual Review of Earth and Planetary Sciences. 17 (1): 169-195. Bibcode:1989AREPS..17..169C. doi:10.1146/annurev.ea.17.050189.001125.

20. ^ Barber, D. J.; Scott, E. R. D. (14 May 2002). "Origin of supposedly biogenic magnetite in the Martian meteorite Allan Hills 84001". Proceedings of the National Academy of Sciences. 99 (10): 6556-6561. Bibcode:2002PNAS...99.6556B. doi:10.1073/pnas.102045799. PMC . PMID 12011420.

21. Wiltschko, Roswitha; Wiltschko, Wolfgang (2014). "Sensing magnetic directions in birds: radical pair processes involving cryptochrome". Biosensors. 4 (3): 221-42. doi:10.3390/bios4030221. PMC . PMID 25587420. Birds can use the geomagnetic field for compass orientation. Behavioral experiments, mostly with migrating passerines, revealed three characteristics of the avian magnetic compass: (1) it works spontaneously only in a narrow functional window around the intensity of the ambient magnetic field, but can adapt to other intensities, (2) it is an "inclination compass", not based on the polarity of the magnetic field, but the axial course of the field lines, and (3) it requires short-wavelength light from UV to 565 nm Green.

22. ^ Kirschvink, Joseph; et al. (1992). "Magnetite biomineralization in the human brain". Proceedings of the National Academy of Sciences of the USA. 89 (16): 7683-7687. Bibcode:1992PNAS...89.7683K. doi:10.1073/pnas.89.16.7683. PMC . PMID 1502184. Using an ultrasensitive superconducting magnetometer in a clean-lab environment, we have detected the presence of ferromagnetic material in a variety of tissues from the human brain.

23. ^ Kirschvink, J L; Kobayashi-Kirschvink, A; Diaz-Ricci, J C; Kirschvink, S J (1992). "Magnetite in human tissues: a mechanism for the biological effects of weak ELF magnetic fields". Bioelectromagnetics. Suppl 1: 101-13. CiteSeerX 10.1.1.326.4179. doi:10.1002/bem.2250130710. PMID 1285705. A simple calculation shows that magnetosomes moving in response to earth-strength ELF fields are capable of opening trans-membrane ion channels, in a fashion similar to those predicted by ionic resonance models. Hence, the presence of trace levels of biogenic magnetite in virtually all human tissues examined suggests that similar biophysical processes may explain a variety of weak field ELF bioeffects.

24. ^ Nakajima, Ken-ichi; Zhu, Kan; Sun, Yao-Hui; Hegyi, Bence; Zeng, Qunli; Murphy, Christopher J; Small, J Victor; Chen-Izu, Ye; Izumiya, Yoshihiro; Penninger, Josef M; Zhao, Min (2015). "KCNJ15/Kir4.2 couples with polyamines to sense weak extracellular electric fields in galvanotaxis". Nature Communications. 6: 8532. Bibcode:2015NatCo...6.8532N. doi:10.1038/ncomms9532. PMC . PMID 26449415. Taken together these data suggest a previously unknown two-molecule sensing mechanism in which KCNJ15/Kir4.2 couples with polyamines in sensing weak electric fields.

25. ^ Lovley, Derek; Stolz, John; Nord, Gordon; Phillips, Elizabeth. "Anaerobic production of magnetite by a dissimilatory iron-reducing microorganism" (PDF). geobacter.org. US Geological Survey, Reston, Virginia 22092, USA Department of Biochemistry, University of Massachusetts, Amherst, Massachusetts 01003, USA. Archived from the original (PDF) on 29 March 2017.

26. ^ Kishkinev, D A; Chernetsov, N S (2014). "[Magnetoreception systems in birds: a review of current research]". Zhurnal Obshcheĭ Biologii. 75 (2): 104-23. PMID 25490840. There are good reasons to believe that this visual magnetoreceptor processes compass magnetic information which is necessary for migratory orientation.

27. ^ Lowenstam, H.A. (1967). "Lepidocrocite, an apatite mineral, and magnetic in teeth of chitons (Polyplacophora)". Science. 156 (3780): 1373-1375. Bibcode:1967Sci...156.1373L. doi:10.1126/science.156.3780.1373. PMID 5610118. S2CID 40567757. X-ray diffraction patterns show that the mature denticles of three extant chiton species are composed of the mineral lepidocrocite and an apatite mineral, probably francolite, in addition to magnetite.

28. ^ Magnetite Nano-Particles in Information Processing: From the Bacteria to the Human Brain Neocortex - ISBN 9781-61761-839-0

29. Barbara A. Maher; Imad A. M. Ahmed; Vassil Karloukovski; Donald A. MacLaren; Penelope G. Foulds; David Allsop; David M. A. Mann; Ricardo Torres-Jardón; Lilian Calderon-Garciduenas (2016). "Magnetite pollution nanoparticles in the human brain". PNAS. 113 (39): 10797-10801. Bibcode:2016PNAS..11310797M. doi:10.1073/pnas.1605941113. PMC . PMID 27601646.

30. ^ Eric Hand (June 23, 2016). "Maverick scientist thinks he has discovered a magnetic sixth sense in humans". Science. doi:10.1126/science.aaf5803.

31. ^ Baker, R R (1988). "Human magnetoreception for navigation". Progress in Clinical and Biological Research. 257: 63-80. PMID 3344279.

32. ^ Kirschvink, Joseph L; Winklhofer, Michael; Walker, Michael M (2010). "Biophysics of magnetic orientation: strengthening the interface between theory and experimental design". Journal of the Royal Society, Interface. 7 Suppl 2 (Suppl 2): S179-91. doi:10.1098/rsif.2009.0491.focus. PMC . PMID 20071390.

33. ^ Maher, B.A.; Ahmed, I.A.; Karloukovski, V.; MacLaren, D.A.; Foulds, P.G.; Allsop, D.; Mann, D.M.; Torres-Jardón, R.; Calderon-Garciduenas, L. (2016). "Magnetite pollution nanoparticles in the human brain". Proceedings of the National Academy of Sciences. 113 (39): 10797-10801. Bibcode:2016PNAS..11310797M. doi:10.1073/pnas.1605941113. PMC . PMID 27601646.

34. ^ Wilson, Clare (5 September 2016). "Air pollution is sending tiny magnetic particles into your brain". New Scientist. 231 (3090).

35. Franz Oeters et al"Iron" in Ullmann's Encyclopedia of Industrial Chemistry, 2006, Wiley-VCH, Weinheim. ^ doi:10.1002/14356007.a14_461.pub2

36. Schoenherr, Steven (2002). "The History of Magnetic Recording". Audio Engineering Society.

37. ^ Jozwiak, W. K.; Kaczmarek, E.; et al. (2007). "Reduction behavior of iron oxides in hydrogen and carbon monoxide atmospheres". Applied Catalysis A: General. 326: 17-27. doi:10.1016/j.apcata.2007.03.021.

38. ^ Appl, Max (2006). "Ammonia". Ullmann's Encyclopedia of Industrial Chemistry. Weinheim: Wiley-VCH. doi:10.1002/14356007.a02_143.pub2.

39. Blaney, Lee (2007). "Magnetite (Fe3O4): Properties, Synthesis, and Applications". The Lehigh Review. 15 (5). Archived from the original on 2020-11-11.

40. ^ Stephen, Zachary R.; Kievit, Forrest M.; Zhang, Miqin (2011). "Magnetite nanoparticles for medical MR imaging". Materials Today. 14 (7-8): 330-338. doi:10.1016/s1369-7021(11)70163-8. PMC . PMID 22389583.

41. ^ Toronto, University of. "Magnetene: Graphene-like 2D material leverages quantum effects to achieve ultra-low friction". phys.org.

# data[].attributes.attachment_link.description
Magnetite is a mineral and one of the main iron ores, with the chemical formula . It is one of the oxides of iron, and is ferrimagnetic; it is attracted to a magnet and can be magnetized to become a permanent magnet itself.

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Non Metallic
Has maroon and grey coloring
Produces maroon streak
Cleavage
Can scratch a penny
Found in ore of iron

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Hematite: The mineral hematite information and pictures

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Detailed description, properties, and locality information guide about the iron mineral hematite.

# data[].attributes.body
Non metallic
Has puprle, white coloring
Produces a clear streak
Fracture
Can scratch a glass plate
Used for glass and jewelry

# data[].attributes.attachment_link.readable_text
Quartz is a hard, crystalline mineral composed of silica ( silicon dioxide). The atoms are linked in a continuous framework of SiO ₄ silicon-oxygen tetrahedra, with each oxygen being shared between two tetrahedra, giving an overall chemical formula of . Quartz is, therefore, classified structurally as a framework silicate mineral and compositionally as an oxide mineral. Quartz is the second most abundant mineral in Earth's continental crust, behind feldspar.

Quartz exists in two forms, the normal α-quartz and the high-temperature β-quartz, both of which are chiral. The transformation from α-quartz to β-quartz takes place abruptly at 573 °C (846 K; 1,063 °F). Since the transformation is accompanied by a significant change in volume, it can easily induce microfracturing of ceramics or rocks passing through this temperature threshold.

There are many different varieties of quartz, several of which are classified as gemstones. Since antiquity, varieties of quartz have been the most commonly used minerals in the making of jewelry and hardstone carvings, especially in Europe and Asia.

Quartz is the mineral defining the value of 7 on the Mohs scale of hardness, a qualitative scratch method for determining the hardness of a material to abrasion.

The word "quartz" is derived from the German word Quarz, which had the same form in the first half of the 14th century in Middle High German and in East Central German and which came from the Polish dialect term , which corresponds to the Czech term tvrdý ("hard"). Some sources, however, attribute the word's origin to the Saxon word Querkluftertz, meaning cross-vein ore.

The Ancient Greeks referred to quartz as κρύσταλλος ( krustallos) derived from the Ancient Greek ( kruos) meaning "icy cold", because some philosophers (including Theophrastus) understood the mineral to be a form of supercooled ice. Today, the term rock crystal is sometimes used as an alternative name for transparent coarsely crystalline quartz.

Roman naturalist Pliny the Elder believed quartz to be water ice, permanently frozen after great lengths of time. He supported this idea by saying that quartz is found near glaciers in the Alps, but not on volcanic mountains, and that large quartz crystals were fashioned into spheres to cool the hands. This idea persisted until at least the 17th century. He also knew of the ability of quartz to split light into a spectrum.

In the 17th century, Nicolas Steno's study of quartz paved the way for modern crystallography. He discovered that regardless of a quartz crystal's size or shape, its long prism faces always joined at a perfect 60° angle.

Quartz belongs to the trigonal crystal system at room temperature, and to the hexagonal crystal system above 573 °C (846 K; 1,063 °F). The ideal crystal shape is a six-sided prism terminating with six-sided pyramid-like rhombohedrons at each end. In nature, quartz crystals are often twinned (with twin right-handed and left-handed quartz crystals), distorted, or so intergrown with adjacent crystals of quartz or other minerals as to only show part of this shape, or to lack obvious crystal faces altogether and appear massive.

Well-formed crystals typically form as a druse (a layer of crystals lining a void), of which quartz geodes are particularly fine examples. The crystals are attached at one end to the enclosing rock, and only one termination pyramid is present. However, doubly terminated crystals do occur where they develop freely without attachment, for instance, within gypsum.

α-quartz crystallizes in the trigonal crystal system, space group P3 ₁21 or P3 ₂21 (space group 152 or 154 resp.) depending on the chirality. Above 573 °C (846 K; 1,063 °F), α-quartz in P3 ₁21 becomes the more symmetric hexagonal P6 ₄22 (space group 181), and α-quartz in P3 ₂21 goes to space group P6 ₂ 22 (no. 180).

These space groups are truly chiral (they each belong to the 11 enantiomorphous pairs). Both α-quartz and β-quartz are examples of chiral crystal structures composed of achiral building blocks (SiO ₄ tetrahedra in the present case). The transformation between α- and β-quartz only involves a comparatively minor rotation of the tetrahedra with respect to one another, without a change in the way they are linked. However, there is a significant change in volume during this transition,[clarification needed ] and this can result in significant microfracturing in ceramics and in rocks of the Earth's crust.

Varieties (according to microstructure)

Although many of the varietal names historically arose from the color of the mineral, current scientific naming schemes refer primarily to the microstructure of the mineral. Color is a secondary identifier for the cryptocrystalline minerals, although it is a primary identifier for the macrocrystalline varieties.

Type

Color and description

Transparency

Rock crystal

Colorless

Transparent

Chalcedony
Fibrous, variously translucent, cryptocrystalline quartz occurring in many varieties.
The term is often used for white, cloudy, or lightly colored material intergrown with moganite.
Otherwise more specific names are used.

Aventurine
Quartz with tiny aligned inclusions (usually mica) that shimmer with aventurescence

Translucent to opaque

Dumortierite quartz

Contains large amounts of blue dumortierite crystals

Translucent

Varieties (according to color)

Quartz crystal demonstrating transparency

Pure quartz, traditionally called rock crystal or clear quartz, is colorless and transparent or translucent and has often been used for hardstone carvings, such as the Lothair Crystal. Common colored varieties include citrine, rose quartz, amethyst, smoky quartz, milky quartz, and others. These color differentiations arise from the presence of impurities which change the molecular orbitals, causing some electronic transitions to take place in the visible spectrum causing colors.

The most important distinction between types of quartz is that of macrocrystalline (individual crystals visible to the unaided eye) and the microcrystalline or cryptocrystalline varieties ( aggregates of crystals visible only under high magnification). The cryptocrystalline varieties are either translucent or mostly opaque, while the transparent varieties tend to be macrocrystalline. Chalcedony is a cryptocrystalline form of silica consisting of fine intergrowths of both quartz, and its monoclinic polymorph moganite. Other opaque gemstone varieties of quartz, or mixed rocks including quartz, often including contrasting bands or patterns of color, are agate, carnelian or sard, onyx, heliotrope, and jasper.

Amethyst is a form of quartz that ranges from a bright vivid violet to a dark or dull lavender shade. The world's largest deposits of amethysts can be found in Brazil, Mexico, Uruguay, Russia, France, Namibia, and Morocco. Sometimes amethyst and citrine are found growing in the same crystal. It is then referred to as ametrine. Amethyst derives its color from traces of iron in its structure.

Blue quartz contains inclusions of fibrous magnesio-riebeckite or crocidolite.

Inclusions of the mineral dumortierite within quartz pieces often result in silky-appearing splotches with a blue hue. Shades of purple or grey sometimes also are present. "Dumortierite quartz" (sometimes called "blue quartz") will sometimes feature contrasting light and dark color zones across the material. "Blue quartz" is a minor gemstone.

Citrine is a variety of quartz whose color ranges from pale yellow to brown due to a submicroscopic distribution of colloidal ferric hydroxide impurities. Natural citrines are rare; most commercial citrines are heat-treated amethysts or smoky quartzes. However, a heat-treated amethyst will have small lines in the crystal, as opposed to a natural citrine's cloudy or smoky appearance. It is nearly impossible to differentiate between cut citrine and yellow topaz visually, but they differ in hardness. Brazil is the leading producer of citrine, with much of its production coming from the state of Rio Grande do Sul. The name is derived from the Latin word citrina which means "yellow" and is also the origin of the word " citron". Sometimes citrine and amethyst can be found together in the same crystal, which is then referred to as ametrine. Citrine has been referred to as the "merchant's stone" or "money stone", due to a superstition that it would bring prosperity.

Citrine was first appreciated as a golden-yellow gemstone in Greece between 300 and 150 BC, during the Hellenistic Age. Yellow quartz was used prior to that to decorate jewelry and tools but it was not highly sought after.

Milk quartz or milky quartz is the most common variety of crystalline quartz. The white color is caused by minute fluid inclusions of gas, liquid, or both, trapped during crystal formation, making it of little value for optical and quality gemstone applications.

Rose quartz is a type of quartz that exhibits a pale pink to rose red hue. The color is usually considered as due to trace amounts of titanium, iron, or manganese in the material. Some rose quartz contains microscopic rutile needles that produce asterism in transmitted light. Recent X-ray diffraction studies suggest that the color is due to thin microscopic fibers of possibly dumortierite within the quartz.

Additionally, there is a rare type of pink quartz (also frequently called crystalline rose quartz) with color that is thought to be caused by trace amounts of phosphate or aluminium. The color in crystals is apparently photosensitive and subject to fading. The first crystals were found in a pegmatite found near Rumford, Maine, US, and in Minas Gerais, Brazil. The crystals found are more transparent and euhedral, due to the impurities of phosphate and aluminium that formed crystalline rose quartz, unlike the iron and microscopic dumortierite fibers that formed rose quartz.

Smoky quartz is a gray, translucent version of quartz. It ranges in clarity from almost complete transparency to a brownish-gray crystal that is almost opaque. Some can also be black. The translucency results from natural irradiation acting on minute traces of aluminum in the crystal structure.

Prase is a green variety of quartz. The green color is caused by inclusions of amphibole.

Prasiolite, also known as vermarine, is a variety of quartz that is green in color. The green is caused by iron ions. It is a rare mineral in nature and is typically found with amethyst; most "prasiolite" is not natural - it has been artificially produced by heating of amethyst. Since 1950 [citation needed ], almost all natural prasiolite has come from a small Brazilian mine, but it is also seen in Lower Silesia in Poland. Naturally occurring prasiolite is also found in the Thunder Bay area of Canada.

Quartz crystals have piezoelectric properties; they develop an electric potential upon the application of mechanical stress. Quartz's piezoelectric properties were discovered by Jacques and Pierre Curie in 1880.

Granite rock in the cliff of Gros la Tête on Aride Island, Seychelles. The thin (1-3 cm wide) brighter layers are quartz veins, formed during the late stages of crystallization of granitic magmas. They are sometimes called "hydrothermal veins".

Quartz is a defining constituent of granite and other felsic igneous rocks. It is very common in sedimentary rocks such as sandstone and shale. It is a common constituent of schist, gneiss, quartzite and other metamorphic rocks. Quartz has the lowest potential for weathering in the Goldich dissolution series and consequently it is very common as a residual mineral in stream sediments and residual soils. Generally a high presence of quartz suggests a " mature" rock, since it indicates the rock has been heavily reworked and quartz was the primary mineral that endured heavy weathering.

While the majority of quartz crystallizes from molten magma, quartz also chemically precipitates from hot hydrothermal veins as gangue, sometimes with ore minerals like gold, silver and copper. Large crystals of quartz are found in magmatic pegmatites. Well-formed crystals may reach several meters in length and weigh hundreds of kilograms.

The largest documented single crystal of quartz was found near Itapore, Goiaz, Brazil; it measured approximately 6.1 m × 1.5 m × 1.5 m and weighed 39,916 kilograms.

Quartz is extracted from open pit mines. Miners occasionally use explosives to expose deep pockets of quartz. More frequently, bulldozers and backhoes are used to remove soil and clay and expose quartz veins, which are then worked using hand tools. Care must be taken to avoid sudden temperature changes that may damage the crystals.

Tridymite and cristobalite are high-temperature polymorphs of SiO ₂ that occur in high-silica volcanic rocks. Coesite is a denser polymorph of SiO ₂ found in some meteorite impact sites and in metamorphic rocks formed at pressures greater than those typical of the Earth's crust. Stishovite is a yet denser and higher-pressure polymorph of SiO ₂ found in some meteorite impact sites. Moganite is a monoclinic polymorph. Lechatelierite is an amorphous silica glass SiO ₂ which is formed by lightning strikes in quartz sand.

As quartz is a form of silica, it is a possible cause for concern in various workplaces. Cutting, grinding, chipping, sanding, drilling, and polishing natural and manufactured stone products can release hazardous levels of very small, crystalline silica dust particles into the air that workers breathe. Crystalline silica of respirable size is a recognized human carcinogen and may lead to other diseases of the lungs such as silicosis and pulmonary fibrosis.

A synthetic quartz crystal grown by the hydrothermal method, about 19 centimetres (7.5 in) long and weighing about 127 grams (4.5 oz)

Not all varieties of quartz are naturally occurring. Some clear quartz crystals can be treated using heat or gamma-irradiation to induce color where it would not otherwise have occurred naturally. Susceptibility to such treatments depends on the location from which the quartz was mined.

Prasiolite, an olive colored material, is produced by heat treatment; natural prasiolite has also been observed in Lower Silesia in Poland. Although citrine occurs naturally, the majority is the result of heat-treating amethyst or smoky quartz. Carnelian has been heat-treated to deepen its color since prehistoric times.

Because natural quartz is often twinned, synthetic quartz is produced for use in industry. Large, flawless, single crystals are synthesized in an autoclave via the hydrothermal process.

Like other crystals, quartz may be coated with metal vapors to give it an attractive sheen.

Quartz is the most common material identified as the mystical substance maban in Australian Aboriginal mythology. It is found regularly in passage tomb cemeteries in Europe in a burial context, such as Newgrange or Carrowmore in Ireland. Quartz was also used in Prehistoric Ireland, as well as many other countries, for stone tools; both vein quartz and rock crystal were knapped as part of the lithic technology of the prehistoric peoples.

While jade has been since earliest times the most prized semi-precious stone for carving in East Asia and Pre-Columbian America, in Europe and the Middle East the different varieties of quartz were the most commonly used for the various types of jewelry and hardstone carving, including engraved gems and cameo gems, rock crystal vases, and extravagant vessels. The tradition continued to produce objects that were very highly valued until the mid-19th century, when it largely fell from fashion except in jewelry. Cameo technique exploits the bands of color in onyx and other varieties.

Efforts to synthesize quartz began in the mid-nineteenth century as scientists attempted to create minerals under laboratory conditions that mimicked the conditions in which the minerals formed in nature: German geologist Karl Emil von Schafhäutl (1803-1890) was the first person to synthesize quartz when in 1845 he created microscopic quartz crystals in a pressure cooker. However, the quality and size of the crystals that were produced by these early efforts were poor.

Elemental impurity incorporation strongly influences the ability to process and utilize quartz. Naturally occurring quartz crystals of extremely high purity, necessary for the crucibles and other equipment used for growing silicon wafers in the semiconductor industry, are expensive and rare. These high-purity quartz are defined as containing less than 50 ppm of impurity elements. A major mining location for high purity quartz is the Spruce Pine Gem Mine in Spruce Pine, North Carolina, United States. Quartz may also be found in Caldoveiro Peak, in Asturias, Spain.

By the 1930s, the electronics industry had become dependent on quartz crystals. The only source of suitable crystals was Brazil; however, World War II disrupted the supplies from Brazil, so nations attempted to synthesize quartz on a commercial scale. German mineralogist Richard Nacken (1884-1971) achieved some success during the 1930s and 1940s. After the war, many laboratories attempted to grow large quartz crystals. In the United States, the U.S. Army Signal Corps contracted with Bell Laboratories and with the Brush Development Company of Cleveland, Ohio to synthesize crystals following Nacken's lead. (Prior to World War II, Brush Development produced piezoelectric crystals for record players.) By 1948, Brush Development had grown crystals that were 1.5 inches (3.8 cm) in diameter, the largest at that time. By the 1950s, hydrothermal synthesis techniques were producing synthetic quartz crystals on an industrial scale, and today virtually all the quartz crystal used in the modern electronics industry is synthetic.

An early use of the piezoelectricity of quartz crystals was in phonograph pickups. One of the most common piezoelectric uses of quartz today is as a crystal oscillator. The quartz oscillator or resonator was first developed by Walter Guyton Cady in 1921. George Washington Pierce designed and patented quartz crystal oscillators in 1923. The quartz clock is a familiar device using the mineral. Warren Marrison created the first quartz oscillator clock based on the work of Cady and Pierce in 1927. The resonant frequency of a quartz crystal oscillator is changed by mechanically loading it, and this principle is used for very accurate measurements of very small mass changes in the quartz crystal microbalance and in thin-film thickness monitors.

Rock crystal jug with cut festoon decoration by Milan workshop from the second half of the 16th century, National Museum in Warsaw. The city of Milan, apart from Prague and Florence, was the main Renaissance centre for crystal cutting.

Synthetic quartz crystals produced in the autoclave shown in Western Electric's pilot hydrothermal quartz plant in 1959

Fatimid ewer in carved rock crystal (clear quartz) with gold lid, c.  1000

Almost all the industrial demand for quartz crystal (used primarily in electronics) is met with synthetic quartz produced by the hydrothermal process. However, synthetic crystals are less prized for use as gemstones. The popularity of crystal healing has increased the demand for natural quartz crystals, which are now often mined in developing countries using primitive mining methods, sometimes involving child labor.

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Quartz is a hard, crystalline mineral composed of silica ( silicon dioxide). The atoms are linked in a continuous framework of SiO 4 silicon-oxygen tetrahedra, with each oxygen being shared between two tetrahedra, giving an overall chemical formula of . Quartz is, therefore, classified structurally as a framework silicate mineral and compositionally as an oxide mineral.

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Has a non metallic luster
It has a dark brown to black coloring
Produces a clear streak
Cleavage
You can break it with your fingernail
Used for electrical purposes

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The uses and properties of the mineral Biotite

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Non metallic
Has a yellow coloring
White streak
Fracture
Can create a scratch with your Fingernail

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Alternative name

Sulphur (British spelling)

Allotropes

see Allotropes of sulfur

Appearance

Lemon yellow sintered microcrystals

(Atomic number Z)

16

Group
group 16 (chalcogens)

Period
period 3

Block

p-block

Electron configuration

[ Ne] 3s ² 3p ⁴

Electrons per shell

2, 8, 6

at Phase STP
solid

Melting point

alpha (α-S ₈): 388.36 K ​(115.21 °C, ​239.38 °F)

(near Density r.t.)

alpha (α-S ₈): 2.07 g/cm ³
beta (β-S ₈): 1.96 g/cm ³
gamma (γ-S ₈): 1.92 g/cm ³

when liquid (at m.p.)

1.819 g/cm ³

Heat of fusion

beta (β-S ₈): 1.727 kJ/mol

Heat of vaporization

beta (β-S ₈): 45 kJ/mol

Vapor pressure

P (Pa)

1

10

100

1 k

10 k

100 k

at T (K)

375

408

449

508

591

717

Oxidation states

−2, −1, 0, +1, +2, +3, +4, +5, +6 (a strongly acidic oxide)

Ionization energies

1st: 999.6 kJ/mol

2nd: 2252 kJ/mol

3rd: 3357 kJ/mol

(more)

Covalent radius

105±3 pm

Spectral lines of sulfur

Natural occurrence
primordial

Crystal structure

alpha (α-S ₈): ​ orthorhombic ( oF128)

Lattice constants

a = 1.0460 nm
b = 1.2861 nm
c = 2.4481 nm (at 20 °C)[3]

Crystal structure

beta (β-S ₈): ​ monoclinic ( mP48)

Lattice constants

a = 1.0923 nm
b = 1.0851 nm
c = 1.0787 nm
β = 95.905° (at 20 °C)[3]

Electrical resistivity

2×10 ¹⁵ Ω⋅m (at 20 °C) (amorphous)

Magnetic ordering

diamagnetic[4]

Molar magnetic susceptibility

alpha (α-S−15.5×10 ₈): −6 cm ³/mol (298 K)[5]

Discovery

before 2000 BCE[6]

Recognized as an element by

Antoine Lavoisier (1777)

³⁴ S abundances vary greatly (between 3.96 and 4.77 percent) in natural samples.

Category: Sulfur
| 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 . 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, which means "burning stone". Today, almost all elemental sulfur is produced as a byproduct of removing sulfur-containing contaminants from natural gas and petroleum. 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.

Sulfur forms several polyatomic molecules. The best-known allotrope is octasulfur, cyclo-S ₈. The point group of cyclo-S ₈ is D 4d and its dipole moment is 0 D. Octasulfur is a soft, bright-yellow solid that is odorless. It melts at 115.21 °C (239.38 °F), boils at 444.6 °C (832.3 °F). At 95.2 °C (203.4 °F), below its melting temperature, cyclo-octasulfur begins slow changing from α-octasulfur to the β- polymorph. The structure of the S ₈ 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), as it predominantly consists of the β-S ₈ 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 200 °C (392 °F). The density of sulfur is about 2 g/cm ³, 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).

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.

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. 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.

The structure of the cyclooctasulfur molecule, S₈

Sulfur forms over 30 solid allotropes, more than any other element. Besides S ₈, several other rings are known. Removing one atom from the crown gives S ₇, which is of a deeper yellow than S ₈. HPLC analysis of "elemental sulfur" reveals an equilibrium mixture of mainly S ₈, but with S ₇ and small amounts of S ₆. Larger rings have been prepared, including S ₁₂ and S ₁₈.

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.

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 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 (δ ³⁴S) 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. However, there are ongoing discussions about what is the real reason of the δ ³⁴ S shifts, biological activity or postdeposital alteration.

For example, when sulfide minerals are precipitated, isotopic equilibration among solids and liquid may cause small differences in the values of co-genetic minerals. The differences between minerals can be used to estimate the temperature of equilibration. The and δ ³⁴S 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 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. 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 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 ³⁴ S values than lakes believed to be dominated by watershed sources of sulfate.

The radioactive ³⁵S is formed in cosmic ray spallation of the atmospheric . 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 + 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 ³⁵ S, its compounds are relatively safe as long as they are not ingested or absorbed by the body.

Most of the yellow and orange hues of Io are due to elemental sulfur and sulfur compounds deposited by active volcanoes.

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 2.5×10 ⁹ 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.

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 cm × 16 cm × 11 cm (8.7 in × 6.3 in × 4.3 in). Historically, Sicily was a major source of sulfur in the Industrial Revolution. 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.

Native sulfur is synthesised 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. 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.

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

SLapis lazuli owes its blue color to a trisulfur radical anion (−
₃ )

Sulfur polycations, , and are produced when sulfur is reacted with oxidising agents in a strongly acidic solution. 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. is deep blue, is yellow and is red.

Reduction of sulfur gives various polysulfides with the formula , 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 gives the blue color of the rock lapis lazuli.

Two parallel sulfur chains grown inside a single-wall carbon nanotube (CNT, a). Zig-zag (b) and straight (c) S chains inside double-wall CNTs[47]

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, H ₂Sx, where x = 2, 3, and 4. Ultimately, reduction of sulfur produces sulfide salts:

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

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

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).

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.

Sulfur reacts with fluorine to give the highly reactive sulfur tetrafluoride and the highly inert sulfur hexafluoride. 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 (SOCl ₂) is a common reagent in organic synthesis. Bromine also oxidizes sulfur to form sulfur dibromide and disulfur dibromide.

Sulfur oxidizes cyanide and sulfite to give thiocyanate and thiosulfate, respectively.

Some of the main classes of sulfur-containing organic compounds include the following:

Compounds with carbon-sulfur multiple bonds are uncommon, an exception being carbon disulfide, a volatile colorless liquid that is structurally similar to carbon dioxide. It is used as a reagent to make the polymer rayon and many organosulfur compounds. Unlike carbon monoxide, carbon monosulfide is stable only as an extremely dilute gas, found between solar systems.

Organosulfur compounds are responsible for some of the unpleasant odors of decaying organic matter. They are widely known as the odorant in domestic natural gas, garlic odor, and skunk spray, as well as a component of bad breath odor. Not all organic sulfur compounds smell unpleasant at all concentrations: the sulfur-containing monoterpenoid grapefruit mercaptan in small concentrations is the characteristic scent of grapefruit, but has a generic thiol odor at larger concentrations. Sulfur mustard, a potent vesicant, was used in World War I as a disabling agent.

Sulfur-sulfur bonds are a structural component used to stiffen rubber, similar to the disulfide bridges that rigidify proteins (see biological below). In the most common type of industrial "curing" or hardening and strengthening of natural rubber, elemental sulfur is heated with the rubber to the point that chemical reactions form disulfide bridges between isoprene units of the polymer. This process, patented in 1843, made rubber a major industrial product, especially in automobile tires. Because of the heat and sulfur, the process was named vulcanization, after the Roman god of the forge and volcanism.

Pharmaceutical container for sulfur from the first half of the 20th century. From the Museo del Objeto del Objeto collection

Being abundantly available in native form, sulfur was known in ancient times and is referred to in the Torah ( Genesis). English translations of the Christian Bible commonly referred to burning sulfur as "brimstone", giving rise to the term " fire-and-brimstone" sermons, in which listeners are reminded of the fate of eternal damnation that await the unbelieving and unrepentant. It is from this part of the Bible that Hell is implied to "smell of sulfur" (likely due to its association with volcanic activity). According to the Ebers Papyrus, a sulfur ointment was used in ancient Egypt to treat granular eyelids. Sulfur was used for fumigation in preclassical Greece; this is mentioned in the . Pliny the Elder discusses sulfur in book 35 of his Natural History, saying that its best-known source is the island of Melos. He mentions its use for fumigation, medicine, and bleaching cloth.

A natural form of sulfur known as shiliuhuang () was known in China since the 6th century BC and found in Hanzhong. By the 3rd century, the Chinese had 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. The Wujing Zongyao of 1044 AD described various formulas for Chinese black powder, which is a mixture of potassium nitrate (), charcoal, and sulfur.

Indian alchemists, practitioners of the "science of chemicals" ( Sanskrit: रसशास्त्र, rasaśāstra), wrote extensively about the use of sulfur in alchemical operations with mercury, from the eighth century AD onwards. In the tradition, sulfur is called "the smelly" (, gandhaka).

Early European alchemists gave sulfur a unique alchemical symbol, a triangle atop a cross (🜍). (This is sometimes confused with the astronomical crossed-spear symbol ⚴ for 2 Pallas.) The variation known as brimstone has a symbol combining a two-barred cross atop a lemniscate (🜏). In traditional skin treatment, elemental sulfur was used (mainly in creams) to alleviate such conditions as scabies, ringworm, psoriasis, eczema, and acne. The mechanism of action is unknown-though elemental sulfur does oxidize slowly to sulfurous acid, which is (through the action of sulfite) a mild reducing and antibacterial agent.

Today sulfur is known to have antifungal, antibacterial, and keratolytic activity; in the past it was used against acne vulgaris, rosacea, seborrheic dermatitis, dandruff, pityriasis versicolor, scabies, and warts.[67] This 1881 advertisement baselessly claims efficacy against rheumatism, gout, baldness, and graying of hair.

Sulfur appears in a column of fixed (non-acidic) alkali in a chemical table of 1718. Antoine Lavoisier used sulfur in combustion experiments, writing of some of these in 1777.

Sulfur deposits in Sicily were the dominant source for more than a century. By the late 18th century, about 2,000 tonnes per year of sulfur were imported into Marseille, France, for the production of sulfuric acid for use in the Leblanc process. In industrializing Britain, with the repeal of tariffs on salt in 1824, demand for sulfur from Sicily surged upward. The increasing British control and exploitation of the mining, refining, and transportation of the sulfur, coupled with the failure of this lucrative export to transform Sicily's backward and impoverished economy, led to the Sulfur Crisis of 1840, when King Ferdinand II gave a monopoly of the sulfur industry to a French firm, violating an earlier 1816 trade agreement with Britain. A peaceful solution was eventually negotiated by France.

In 1867, elemental sulfur was discovered in underground deposits in Louisiana and Texas. The highly successful Frasch process was developed to extract this resource.

With the advent of the contact process, the majority of sulfur today is used to make sulfuric acid for a wide range of uses, particularly fertilizer.

In recent times, the main source of sulfur has become petroleum and natural gas. This is due to the requirement to remove sulfur from fuels in order to prevent acid rain, and has resulted in a surplus of sulfur.

Sulfur is derived from the Latin word , which was Hellenized to in the erroneous belief that the Latin word came from Greek. This spelling was later reinterpreted as representing an /f/ sound and resulted in the spelling , which appears in Latin toward the end of the Classical period. The true Ancient Greek word for sulfur, , theîon (from earlier , théeion), is the source of the international chemical prefix . The Modern Standard Greek word for sulfur is θείο, theío.

In 12th-century Anglo-French, it was . In the 14th century, the erroneously Hellenized Latin was restored in Middle English . By the 15th century, both full Latin spelling variants sulfur and sulphur became common in English. The parallel f~ph spellings continued in Britain until the 19th century, when the word was standardized as sulphur. On the other hand, sulfur was the form chosen in the United States, whereas Canada uses both.

The IUPAC adopted the spelling sulfur in 1990 or 1971, depending on the source cited, as did the Nomenclature Committee of the Royal Society of Chemistry in 1992, restoring the spelling sulfur to Britain. Oxford Dictionaries note that "in chemistry and other technical uses ... the -f- spelling is now the standard form for this and related words in British as well as US contexts, and is increasingly used in general contexts as well."

Traditional sulfur mining at Ijen Volcano, East Java, Indonesia. This image shows the dangerous and rugged conditions the miners face, including toxic smoke and high drops, as well as their lack of protective equipment. The pipes over which they are standing are for condensing sulfur vapors.

Sulfur may be found by itself and historically was usually obtained in this form; pyrite has also been a source of sulfur. In volcanic regions in Sicily, in ancient times, it was found on the surface of the Earth, and the " Sicilian process" was used: sulfur deposits were piled and stacked in brick kilns built on sloping hillsides, with airspaces between them. Then, some sulfur was pulverized, spread over the stacked ore and ignited, causing the free sulfur to melt down the hills. Eventually the surface-borne deposits played out, and miners excavated veins that ultimately dotted the Sicilian landscape with labyrinthine mines. Mining was unmechanized and labor-intensive, with pickmen freeing the ore from the rock, and mine-boys or carrying baskets of ore to the surface, often through a mile or more of tunnels. Once the ore was at the surface, it was reduced and extracted in smelting ovens. The conditions in Sicilian sulfur mines were horrific, prompting Booker T. Washington to write "I am not prepared just now to say to what extent I believe in a physical hell in the next world, but a sulfur mine in Sicily is about the nearest thing to hell that I expect to see in this life." Sulfur is still mined from surface deposits in poorer nations with volcanoes, such as Indonesia, and worker conditions have not improved much since Booker T. Washington's days.

Elemental sulfur was extracted from salt domes (in which it sometimes occurs in nearly pure form) until the late 20th century. Sulfur is now produced as a side product of other industrial processes such as in oil refining, in which sulfur is undesired. As a mineral, native sulfur under salt domes is thought to be a fossil mineral resource, produced by the action of anaerobic bacteria on sulfate deposits. It was removed from such salt-dome mines mainly by the Frasch process. In this method, superheated water was pumped into a native sulfur deposit to melt the sulfur, and then compressed air returned the 99.5% pure melted product to the surface. Throughout the 20th century this procedure produced elemental sulfur that required no further purification. Due to a limited number of such sulfur deposits and the high cost of working them, this process for mining sulfur has not been employed in a major way anywhere in the world since 2002.

Sulfur recovered from hydrocarbons in Alberta, stockpiled for shipment in North Vancouver, British Columbia

Today, sulfur is produced from petroleum, natural gas, and related fossil resources, from which it is obtained mainly as hydrogen sulfide. Organosulfur compounds, undesirable impurities in petroleum, may be upgraded by subjecting them to hydrodesulfurization, which cleaves the C-S bonds:

Owing to the high sulfur content of the Athabasca Oil Sands, stockpiles of elemental sulfur from this process now exist throughout Alberta, Canada. Another way of storing sulfur is as a binder for concrete, the resulting product having some desirable properties (see sulfur concrete).

The world production of sulfur in 2011 amounted to 69 million tonnes (Mt), with more than 15 countries contributing more than 1 Mt each. Countries producing more than 5 Mt are China (9.6), the United States (8.8), Canada (7.1) and Russia (7.1). Production has been slowly increasing from 1900 to 2010; the price was unstable in the 1980s and around 2010.

Elemental sulfur is used mainly as a precursor to other chemicals. Approximately 85% (1989) is converted to sulfuric acid (H ₂SO ₄):

In 2010, the United States produced more sulfuric acid than any other inorganic industrial chemical. The principal use for the acid is the extraction of phosphate ores for the production of fertilizer manufacturing. Other applications of sulfuric acid include oil refining, wastewater processing, and mineral extraction.

Sulfur reacts directly with methane to give carbon disulfide, which is used to manufacture cellophane and rayon. One of the uses of elemental sulfur is in vulcanization of rubber, where polysulfide chains crosslink organic polymers. Large quantities of sulfites are used to bleach paper and to preserve dried fruit. Many surfactants and detergents (e.g. sodium lauryl sulfate) are sulfate derivatives. Calcium sulfate, gypsum (CaSO ₄·2H ₂O) is mined on the scale of 100 million tonnes each year for use in Portland cement and fertilizers.

When silver-based photography was widespread, sodium and ammonium thiosulfate were widely used as "fixing agents". Sulfur is a component of gunpowder ("black powder").

Amino acids synthesized by living organisms such as methionine and cysteine contain organosulfur groups ( thioester and thiol respectively). The antioxidant glutathione protecting many living organisms against free radicals and oxidative stress also contains organic sulfur. Some crops such as onion and garlic also produce different organosulfur compounds such as syn-propanethial-S-oxide responsible of lacrymal irritation (onions), or diallyl disulfide and allicin (garlic). Sulfates, commonly found in soils and groundwaters are often a sufficient natural source of sulfur for plants and bacteria. Atmospheric deposition of sulfur dioxide (SO ₂) is also a common artificial source ( coal combustion) of sulfur for the soils. Under normal circumstances, in most agricultural soils, sulfur is not a limiting nutrient for plants and microorganisms (see Liebig's barrel). However, in some circumstances, soils can be depleted in sulfate, e.g. if this later is leached by meteoric water ( rain) or if the requirements in sulfur for some types of crops are high. This explains that sulfur is increasingly recognized and used as a component of fertilizers. The most important form of sulfur for fertilizer is calcium sulfate, commonly found in nature as the mineral gypsum (CaSO ₄·2H ₂O). Elemental sulfur is hydrophobic (not soluble in water) and cannot be used directly by plants. Elemental sulfur (ES) is sometimes mixed with bentonite to amend depleted soils for crops with high requirement in organo-sulfur. Over time, oxidation abiotic processes with atmospheric oxygen and soil bacteria can oxidize and convert elemental sulfur to soluble derivatives, which can then be used by microorganisms and plants. Sulfur improves the efficiency of other essential plant nutrients, particularly nitrogen and phosphorus. Biologically produced sulfur particles are naturally hydrophilic due to a biopolymer coating and are easier to disperse over the land in a spray of diluted slurry, resulting in a faster uptake by plants.

The plants requirement for sulfur equals or exceeds the requirement for phosphorus. It is an essential nutrient for plant growth, root nodule formation of legumes, and immunity and defense systems. Sulfur deficiency has become widespread in many countries in Europe. Because atmospheric inputs of sulfur continue to decrease, the deficit in the sulfur input/output is likely to increase unless sulfur fertilizers are used. Atmospheric inputs of sulfur decrease because of actions taken to limit acid rains.

Elemental sulfur is one of the oldest fungicides and pesticides. "Dusting sulfur", elemental sulfur in powdered form, is a common fungicide for grapes, strawberry, many vegetables and several other crops. It has a good efficacy against a wide range of powdery mildew diseases as well as black spot. In organic production, sulfur is the most important fungicide. It is the only fungicide used in organically farmed apple production against the main disease apple scab under colder conditions. Biosulfur (biologically produced elemental sulfur with hydrophilic characteristics) can also be used for these applications.

Standard-formulation dusting sulfur is applied to crops with a sulfur duster or from a dusting plane. Wettable sulfur is the commercial name for dusting sulfur formulated with additional ingredients to make it water miscible. It has similar applications and is used as a fungicide against mildew and other mold-related problems with plants and soil.

Elemental sulfur powder is used as an " organic" (i.e., "green") insecticide (actually an acaricide) against ticks and mites. A common method of application is dusting the clothing or limbs with sulfur powder.

A diluted solution of lime sulfur (made by combining calcium hydroxide with elemental sulfur in water) is used as a dip for pets to destroy ringworm (fungus), mange, and other dermatoses and parasites.

Sulfur candles of almost pure sulfur were burned to fumigate structures and wine barrels, but are now considered too toxic for residences.

Sulfur (specifically octasulfur, S ₈) is used in pharmaceutical skin preparations for the treatment of acne and other conditions. It acts as a keratolytic agent and also kills bacteria, fungi, scabies mites, and other parasites. Precipitated sulfur and colloidal sulfur are used, in form of lotions, creams, powders, soaps, and bath additives, for the treatment of acne vulgaris, acne rosacea, and seborrhoeic dermatitis.

Many drugs contain sulfur. Early examples include antibacterial sulfonamides, known as sulfa drugs. A more recent example is mucolytic acetylcysteine. Sulfur is a part of many bacterial defense molecules. Most β-lactam antibiotics, including the penicillins, cephalosporins and monobactams contain sulfur.

Due to their high energy density and the availability of sulfur, there is ongoing research in creating rechargeable lithium-sulfur batteries. Until now, carbonate electrolytes have caused failures in such batteries after a single cycle. In February 2022, researchers at Drexel University have not only created a prototypical battery that lasted 4000 recharge cycles, but also found the first monoclinic gamma sulfur that remained stable below 95 degrees Celsius.

Sulfur is an essential component of all living cells. It is the eighth most abundant element in the human body by weight, about equal in abundance to potassium, and slightly greater than sodium and chlorine. A 70 kg (150 lb) human body contains about 140 grams (4.9 oz) of sulfur. The main dietary source of sulfur for humans is sulfur-containing amino-acids, which can be found in plant and animal proteins.

In the 1880s, while studying (a bacterium living in a sulfur rich environment), Sergei Winogradsky found that it oxidized hydrogen sulfide (H ₂S) as an energy source, forming intracellular sulfur droplets. Winogradsky referred to this form of metabolism as inorgoxidation (oxidation of inorganic compounds). Another contributor, who continued to study it was Selman Waksman. Primitive bacteria that live around deep ocean volcanic vents oxidize hydrogen sulfide for their nutrition, as discovered by Robert Ballard.

Sulfur oxidizers can use as energy sources reduced sulfur compounds, including hydrogen sulfide, elemental sulfur, sulfite, thiosulfate, and various polythionates (e.g., tetrathionate). They depend on enzymes such as sulfur oxygenase and sulfite oxidase to oxidize sulfur to sulfate. Some lithotrophs can even use the energy contained in sulfur compounds to produce sugars, a process known as chemosynthesis. Some bacteria and archaea use hydrogen sulfide in place of water as the electron donor in chemosynthesis, a process similar to photosynthesis that produces sugars and uses oxygen as the electron acceptor. Sulfur-based chemosynthesis may be simplifiedly compared with photosynthesis:

There are bacteria combining these two ways of nutrition: green sulfur bacteria and purple sulfur bacteria. Also sulfur-oxidizing bacteria can go into symbiosis with larger organisms, enabling the later to use hydrogen sulfide as food to be oxidized. Example: the giant tube worm.

There are sulfate-reducing bacteria, that, by contrast, "breathe sulfate" instead of oxygen. They use organic compounds or molecular hydrogen as the energy source. They use sulfur as the electron acceptor, and reduce various oxidized sulfur compounds back into sulfide, often into hydrogen sulfide. They can grow on other partially oxidized sulfur compounds (e.g. thiosulfates, thionates, polysulfides, sulfites).

There are studies pointing that many deposits of native sulfur in places that were the bottom of the ancient oceans have biological origin. These studies indicate that this native sulfur have been obtained through biological activity, but what is responsible for that (sulfur-oxidizing bacteria or sulfate-reducing bacteria) is still unknown for sure.

Sulfur is absorbed by plants roots from soil as sulfate and transported as a phosphate ester. Sulfate is reduced to sulfide via sulfite before it is incorporated into cysteine and other organosulfur compounds.

→ → → cysteine (thiol) → methionine (thioether)

While the plants' role in transferring sulfur to animals by food chains is more or less understood, the role of sulfur bacteria is just getting investigated.

In all forms of life, most of the sulfur is contained in two proteinogenic amino acids ( cysteine and methionine), thus the element is present in all proteins that contain these amino acids, as well as in respective peptides. Some of the sulfur is comprised in certain metabolites-many of which are cofactors-and sulfated polysaccharides of connective tissue ( chondroitin sulfates, heparin).

Proteins, to execute their biological function, need to have specific space geometry. Formation of this geometry is performed in a process called protein folding, and is provided by intra- and inter-molecular bonds. The process has several stages. While at premier stages a polypeptide chain folds due to hydrogen bonds, at later stages folding is provided (apart from hydrogen bonds) by covalent bonds between two sulfur atoms of two cysteine residues (so called disulfide bridges) at different places of a chain (tertiary protein structure) as well as between two cysteine residues in two separated protein subunits (quaternary protein structure). Both structures easily may be seen in insulin. As the bond energy of a covalent disulfide bridge is higher than the energy of a coordinate bond or hydrophobic interaction, higher disulfide bridges content leads to higher energy needed for protein denaturation. In general disulfide bonds are necessary in proteins functioning outside cellular space, and they do not change proteins' conformation (geometry), but serve as its stabilizers. Within cytoplasm cysteine residues of proteins are saved in reduced state (i.e. in -SH form) by thioredoxins.

This property manifests in following examples. Lysozyme is stable enough to be applied as a drug. Feathers and hair have relative strength, and consisting in them keratin is considered indigestible by most organisms. However, there are fungi and bacteria containing keratinase, and are able to destruct keratin.

Many important cellular enzymes use prosthetic groups ending with -SH moieties to handle reactions involving acyl-containing biochemicals: two common examples from basic metabolism are coenzyme A and alpha-lipoic acid. Cysteine-related metabolites homocysteine and taurine are other sulfur-containing amino acids that are similar in structure, but not coded by DNA, and are not part of the primary structure of proteins, take part in various locations of mammalian physiology. Two of the 13 classical vitamins, biotin and thiamine, contain sulfur, and serve as cofactors to several enzymes. In intracellular chemistry, sulfur operates as a carrier of reducing hydrogen and its electrons for cellular repair of oxidation. Reduced glutathione, a sulfur-containing tripeptide, is a reducing agent through its sulfhydryl (-SH) moiety derived from cysteine.

Methanogenesis, the route to most of the world's methane, is a multistep biochemical transformation of carbon dioxide. This conversion requires several organosulfur cofactors. These include coenzyme M, , the immediate precursor to methane.

Metalloproteins-in which the active site is a transition metal ion (or metal-sulfide cluster) often coordinated by sulfur atoms of cysteine residues-are essential components of enzymes involved in electron transfer processes. Examples include plastocyanin (Cu ²⁺) and nitrous oxide reductase (Cu-S). The function of these enzymes is dependent on the fact that the transition metal ion can undergo redox reactions. Other examples include many zinc proteins, as well as iron-sulfur clusters. Most pervasive are the ferrodoxins, which serve as electron shuttles in cells. In bacteria, the important nitrogenase enzymes contain an Fe-Mo-S cluster and is a catalyst that performs the important function of nitrogen fixation, converting atmospheric nitrogen to ammonia that can be used by microorganisms and plants to make proteins, DNA, RNA, alkaloids, and the other organic nitrogen compounds necessary for life.

Sulfur is also present in molybdenum cofactor.

Easiness of electron flow in a cluster provides catalytic effect of a respective enzyme.

In humans methionine is an essential amino acid; cysteine is conditionally essential and may be synthesized from non-essential serine (sulfur donor would be methionine in this case). Dietary deficiency rarely happens in common conditions. Artificial methionine deficiency is attempted to apply in cancer treatment, but the method is still potentially dangerous.

Isolated sulfite oxidase deficiency is a rare, fatal genetic disease preventing production of sulfite oxidase, needed to metabolize sulfites to sulfates.

Though elemental sulfur is only minimally absorbed through the skin and is of low toxicity to humans, inhalation of sulfur dust or contact with eyes or skin may cause irritation. Excessive ingestion of sulfur can cause a burning sensation or diarrhea, and cases of life-threatening metabolic acidosis have been reported after patients deliberately consumed sulfur as a folk remedy.

When sulfur burns in air, it produces sulfur dioxide. In water, this gas produces sulfurous acid and sulfites; sulfites are antioxidants that inhibit growth of aerobic bacteria and a useful food additive in small amounts. At high concentrations these acids harm the lungs, eyes, or other tissues. In organisms without lungs such as insects, sulfite in high concentration prevents respiration.

Sulfur trioxide (made by catalysis from sulfur dioxide) and sulfuric acid are similarly highly acidic and corrosive in the presence of water. Concentrated sulfuric acid is a strong dehydrating agent that can strip available water molecules and water components from sugar and organic tissue.

The burning of coal and/or petroleum by industry and power plants generates sulfur dioxide (SO ₂) that reacts with atmospheric water and oxygen to produce sulfurous acid (H ₂SO ₃). These acids are components of acid rain, lowering the pH of soil and freshwater bodies, sometimes resulting in substantial damage to the environment and chemical weathering of statues and structures. Fuel standards increasingly require that fuel producers extract sulfur from fossil fuels to prevent acid rain formation. This extracted and refined sulfur represents a large portion of sulfur production. In coal-fired power plants, flue gases are sometimes purified. More modern power plants that use synthesis gas extract the sulfur before they burn the gas.

Hydrogen sulfide is about one-half as toxic as hydrogen cyanide, and intoxicates by the same mechanism (inhibition of the respiratory enzyme cytochrome oxidase), though hydrogen sulfide is less likely to cause sudden poisonings from small inhaled amounts (near its permissible exposure limit (PEL) of 20 ppm) because of its disagreeable odor. However, its presence in ambient air at concentration over 100-150 ppm quickly deadens the sense of smell, and a victim may breathe increasing quantities without noticing until severe symptoms cause death. Dissolved sulfide and hydrosulfide salts are toxic by the same mechanism.

Sigel, Astrid; Freisinger, Eva; Sigel, Roland K.O., eds. (2020). Transition Metals and Sulfur: A Strong Relationship for Life. Guest Editors Martha E Sosa Torres and Peter M.H.Kroneck. Berlin/Boston: de Gruyter. pp. xlv+455. ISBN 978-3-11-058889-7.

# data[].attributes.attachment_link.description
Alternative name Sulphur (British spelling) Allotropes see Allotropes of sulfur Appearance Lemon yellow sintered microcrystals (Atomic number Z) 16 Group group 16 (chalcogens) Period period 3 Block p-block Electron configuration [ Ne] 3s 2 3p 4 Electrons per shell 2, 8, 6 at Phase STP solid Melting point alpha (α-S 8): 388.36 K ​(115.21 °C, ​239.38 °F) (near Density r.t.)

# data[].attributes.body
Metallic luster
Dark gray color
gray streak
Fracture
You can break it with your fingernail
Found in pencil lead

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Graphite ( ), archaically referred to as plumbago, is a crystalline allotrope of carbon, a semimetal, a native element mineral, and a form of coal. Graphite is the most stable form of carbon under standard conditions. Therefore, it is used in thermochemistry as the standard state for defining the heat of formation of carbon compounds.

# data[].attributes.body
Has metallic luster
Has a silver coloring
I has a gray streak
Fracture
You can create a scrath with your fingernail
Used in the ore of lead iron

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Galena, also called lead glance, is the natural mineral form of lead(II) sulfide (PbS). It is the most important ore of lead and an important source of silver.

Galena is one of the most abundant and widely distributed sulfide minerals. It crystallizes in the cubic crystal system often showing octahedral forms. It is often associated with the minerals sphalerite, calcite and fluorite.

Galena is the main ore of lead, used since ancient times, since lead can be smelted from galena in an ordinary wood fire. Galena typically is found in hydrothermal veins in association with sphalerite, marcasite, chalcopyrite, cerussite, anglesite, dolomite, calcite, quartz, barite, and fluorite. It is also found in association with sphalerite in low-temperature lead- zinc deposits within limestone beds. Minor amounts are found in contact metamorphic zones, in pegmatites, and disseminated in sedimentary rock.

In some deposits the galena contains up to 0.5% silver, a byproduct that far outweighs the main lead ore in revenue. In these deposits significant amounts of silver occur as included silver sulfide mineral phases or as limited silver in solid solution within the galena structure. These argentiferous galenas have long been an important ore of silver. Silver-bearing galena is almost entirely of hydrothermal origin; galena in lead-zinc deposits contains little silver.

Galena deposits are found worldwide in various environments. Noted deposits include those at Freiberg in Saxony; Cornwall, the Mendips in Somerset, Derbyshire, and Cumberland in England; the Madan and Rhodope Mountains in Bulgaria; the Sullivan Mine of British Columbia; Broken Hill and Mount Isa in Australia; and the ancient mines of Sardinia.

In the United States, it occurs most notably as lead-zinc ore in the Mississippi Valley type deposits of the Lead Belt in southeastern Missouri, which is the largest known deposit, and in the Driftless Area of Illinois, Iowa and Wisconsin. Galena also was a major mineral of the zinc-lead mines of the tri-state district around Joplin in southwestern Missouri and the adjoining areas of Kansas and Oklahoma. Galena is also an important ore mineral in the silver mining regions of Colorado, Idaho, Utah and Montana. Of the latter, the Coeur d'Alene district of northern Idaho was most prominent.

Australia is world's leading producer of lead as of 2021, most of which is extracted as galena. Argentiferous galena was accidentally discovered at Glen Osmond in 1841, and additional deposits were discovered near Broken Hill in 1876 and at Mount Isa in 1923. Most galena in Australia is found in hydrothermal deposits emplaced around 1680 million years ago, which have since been heavily metamorphosed.

The largest documented crystal of galena is composite cubo-octahedra from the Great Laxey Mine, Isle of Man, measuring 25 cm × 25 cm × 25 cm (10 in × 10 in × 10 in).

Galena is the official state mineral of the U.S. states of Kansas, Missouri, and Wisconsin; the former mining communities of Galena, Kansas, and Galena, Illinois, take their names from deposits of this mineral.

Galena belongs to the octahedral sulfide group of minerals that have metal ions in octahedral positions, such as the iron sulfide pyrrhotite and the nickel arsenide niccolite. The galena group is named after its most common member, with other isometric members that include manganese bearing alabandite and niningerite.

Divalent lead (Pb) cations and sulfur (S) anions form a close-packed cubic unit cell much like the mineral halite of the halide mineral group. Zinc, cadmium, iron, copper, antimony, arsenic, bismuth and selenium also occur in variable amounts in galena. Selenium substitutes for sulfur in the structure constituting a solid solution series. The lead telluride mineral altaite has the same crystal structure as galena.

Within the weathering or oxidation zone galena alters to anglesite (lead sulfate) or cerussite (lead carbonate). Galena exposed to acid mine drainage can be oxidized to anglesite by naturally occurring bacteria and archaea, in a process similar to bioleaching.

One of the oldest uses of galena was in the eye cosmetic kohl. In Ancient Egypt, this was applied around the eyes to reduce the glare of the desert sun and to repel flies, which were a potential source of disease.

In pre-Columbian North America, galena was used by indigenous peoples as an ingredient in decorative paints and cosmetics, and widely traded throughout the eastern United States. Traces of galena are frequently found at the Mississippian city at Kincaid Mounds in present-day Illinois. The galena used at the site originated from deposits in southeastern and central Missouri and the Upper Mississippi Valley.

Galena is the primary ore of lead, and is often mined for its silver content.

It can be used as a source of lead in ceramic glaze.

Galena is a semiconductor with a small band gap of about 0.4 eV, which found use in early wireless communication systems. It was used as the crystal in crystal radio receivers, in which it was used as a point-contact diode capable of rectifying alternating current to detect the radio signals. The galena crystal was used with a sharp wire, known as a "cat's whisker" in contact with it.

1. ^ Anthony, John W.; Bideaux, Richard A.; Bladh, Kenneth W.; Nichols, Monte C., eds. (1990). "Galena". Handbook of Mineralogy (PDF). 1. Chantilly, VA: Mineralogical Society of America. ISBN 0962209708.

2. ^ Young, Courtney A.; Taylor, Patrick R.; Anderson, Corby G. (2008). Hydrometallurgy 2008: Proceedings of the Sixth International Symposium. SME. ISBN 9780873352666.

3. ^ Lucas, A. (May 1928). "Silver in Ancient Times". The Journal of Egyptian Archaeology. 14 (1): 313-319. doi:10.1177/030751332801400160.

4. ^ Klein, Cornelis; Hurlbut, Cornelius S., Jr. (1993). Manual of mineralogy : (after James D. Dana) (21st ed.). New York: Wiley. pp. 354-356. ISBN 047157452X.

5. ^ Walters, Stephen; Bailey, Andrew (1998-12-01). "Geology and mineralization of the Cannington Ag-Pb-Zn deposit; an example of Broken Hill-type mineralization in the eastern succession, Mount Isa Inlier, Australia". Economic Geology. 93 (8): 1307-1329. doi:10.2113/gsecongeo.93.8.1307.

6. ^ "Office of the Secretary of State, Missouri - State Symbols". State of Missouri.

7. ^ "Wisconsin State Symbols". State of Wisconsin. Archived from the original on 2010-01-12.

8. Kansas Place-Names, John Rydjord, University of Oklahoma Press, 1972, p. 77 ^ ISBN 0-8061-0994-7

9. ^ Gannett, Henry (1905). The Origin of Certain Place Names in the United States. Govt. Print. Off. p. 133.

10. ^ Da Silva, Gabriel (2004). "Kinetics and mechanism of the bacterial and ferric sulphate oxidation of galena". Hydrometallurgy. 75: 99. doi:10.1016/j.hydromet.2004.07.001.

11. ^ "Lead pollution from Native Americans attributed to crushing galena for glitter paint, adornments". Indiana University-Purdue University Indianapolis. 21 October 2019.

12. ^ The Glittery Legacy of Lead at a Historic Native American Site, Atlas Obscura, November 7, 2019

# data[].attributes.attachment_link.description
Galena, also called lead glance, is the natural mineral form of lead(II) sulfide (PbS). It is the most important ore of lead and an important source of silver. Galena is one of the most abundant and widely distributed sulfide minerals.

# data[].attributes.body
Has a metallic luster
Has brassy yellow coloring
Produces a colorless streak
Fracture
It takes the strenght of a glass plate in order to create a scratch
Used in the ore of sulfer

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The mineral pyrite (), or iron pyrite, also known as fool's gold, is an iron sulfide with the chemical formula Fe S ₂ (iron (II) disulfide). Pyrite is the most abundant sulfide mineral.

Pyrite's metallic luster and pale brass-yellow hue give it a superficial resemblance to gold, hence the well-known nickname of fool's gold. The color has also led to the nicknames brass, brazzle, and Brazil, primarily used to refer to pyrite found in coal.

The name pyrite is derived from the Greek πυρίτης λίθος ( pyritēs lithos), "stone or mineral which strikes fire", in turn from πῦρ ( pyr), "fire". In ancient Roman times, this name was applied to several types of stone that would create sparks when struck against steel; Pliny the Elder described one of them as being brassy, almost certainly a reference to what we now call pyrite.

By Georgius Agricola's time, c. 1550, the term had become a generic term for all of the sulfide minerals.

Pyrite is usually found associated with other sulfides or oxides in quartz veins, sedimentary rock, and metamorphic rock, as well as in coal beds and as a replacement mineral in fossils, but has also been identified in the sclerites of scaly-foot gastropods. Despite being nicknamed fool's gold, pyrite is sometimes found in association with small quantities of gold. A substantial proportion of the gold is "invisible gold" incorporated into the pyrite (see Carlin-type gold deposit). It has been suggested that the presence of both gold and arsenic is a case of coupled substitution but as of 1997 the chemical state of the gold remained controversial.

Pyrite enjoyed brief popularity in the 16th and 17th centuries as a source of ignition in early firearms, most notably the wheellock, where a sample of pyrite was placed against a circular file to strike the sparks needed to fire the gun.

Pyrite has been used since classical times to manufacture copperas ( iron(II) sulfate). Iron pyrite was heaped up and allowed to weather (an example of an early form of heap leaching). The acidic runoff from the heap was then boiled with iron to produce iron sulfate. In the 15th century, new methods of such leaching began to replace the burning of sulfur as a source of sulfuric acid. By the 19th century, it had become the dominant method.

Pyrite remains in commercial use for the production of sulfur dioxide, for use in such applications as the paper industry, and in the manufacture of sulfuric acid. Thermal decomposition of pyrite into FeS ( iron(II) sulfide) and elemental sulfur starts at 540 °C (1,004 °F); at around 700 °C (1,292 °F), p is about 1 atm.

A newer commercial use for pyrite is as the cathode material in Energizer brand non-rechargeable lithium batteries.

Pyrite is a semiconductor material with a band gap of 0.95 eV. Pure pyrite is naturally n-type, in both crystal and thin-film forms, potentially due to sulfur vacancies in the pyrite crystal structure acting as n-dopants.

During the early years of the 20th century, pyrite was used as a mineral detector in radio receivers, and is still used by crystal radio hobbyists. Until the vacuum tube matured, the crystal detector was the most sensitive and dependable detector available-with considerable variation between mineral types and even individual samples within a particular type of mineral. Pyrite detectors occupied a midway point between galena detectors and the more mechanically complicated perikon mineral pairs. Pyrite detectors can be as sensitive as a modern 1N34A germanium diode detector.

Pyrite has been proposed as an abundant, non-toxic, inexpensive material in low-cost photovoltaic solar panels. Synthetic iron sulfide was used with copper sulfide to create the photovoltaic material. More recent efforts are working toward thin-film solar cells made entirely of pyrite.

Pyrite is used to make marcasite jewelry. Marcasite jewelry, made from small faceted pieces of pyrite, often set in silver, was known since ancient times and was popular in the Victorian era. At the time when the term became common in jewelry making, "marcasite" referred to all iron sulfides including pyrite, and not to the orthorhombic FeS ₂ mineral marcasite which is lighter in color, brittle and chemically unstable, and thus not suitable for jewelry making. Marcasite jewelry does not actually contain the mineral marcasite. The specimens of pyrite, when it appears as good quality crystals, are used in decoration. They are also very popular in mineral collecting. Among the sites that provide the best specimens, highlights the exploited in Navajún, La Rioja (Spain).

China represents the main importing country with an import of around 376,000 tonnes, which resulted at 45% of total global imports. China is also the fastest growing in terms of the unroasted iron pyrites imports, with a CAGR of +27.8% from 2007 to 2016. In value terms, China ($47 million) constitutes the largest market for imported unroasted iron pyrites worldwide, making up 65% of global imports.

In July 2020 scientists reported that they have transformed the abundant diamagnetic material into a ferromagnetic one by inducing voltage, which may lead to applications in devices such as solar cells or magnetic data storage. Researchers at Trinity University Dublin, Ireland have demonstrated that FeS ₂ can be exfoliated into few-layers just like other two-dimensional layered materials such as graphene by a simple liquid-phase exfoliation route. This is the first study to demonstrate the production of non-layered 2D-platelets from 3D bulk FeS ₂. Furthermore, they have used these 2D-platelets with 20% single walled carbon-nanotube as an anode material in lithium-ion batteries, reaching a capacity of 1000 mAh/g close to the theoretical capacity of FeS ₂.

Formal oxidation states for pyrite, marcasite, and arsenopyrite

From the perspective of classical inorganic chemistry, which assigns formal oxidation states to each atom, pyrite is probably best described as Fe ²⁺S ₂ 2−. This formalism recognizes that the sulfur atoms in pyrite occur in pairs with clear S-S bonds. These persulfide units can be viewed as derived from hydrogen disulfide, H ₂S ₂. Thus pyrite would be more descriptively called iron persulfide, not iron disulfide. In contrast, molybdenite, MoS ₂, features isolated sulfide (S 2−) centers and the oxidation state of molybdenum is Mo ⁴⁺. The mineral arsenopyrite has the formula Fe AsS. Whereas pyrite has S ₂ subunits, arsenopyrite has [AsS] units, formally derived from deprotonation of H ₂AsSH. Analysis of classical oxidation states would recommend the description of arsenopyrite as Fe ³⁺[AsS] 3−.

Iron-pyrite FeSS ₂ represents the prototype compound of the crystallographic pyrite structure. The structure is simple cubic and was among the first crystal structures solved by X-ray diffraction.[31] It belongs to the crystallographic space group Pa3 and is denoted by the Strukturbericht notation C2. Under thermodynamic standard conditions the lattice constant of stoichiometric iron pyrite FeS ₂ amounts to 541.87 pm.[32] The unit cell is composed of a Fe face-centered cubic sublattice into which the simple cubic with four iron atoms per cubic unit cell, not face-centered cubic. The iron atoms in the faces are not equivalent by translation alone to the iron atoms at the corners.) The pyrite structure is also used by other compounds
₂ ions are embedded. (Note though that the structure is MX ₂ of transition metals M and chalcogens X = O, S, Se and Te. Also certain dipnictides with X standing for P, As and Sb etc. are known to adopt the pyrite structure.[33]

In the first bonding sphere, the Fe atoms are surrounded by six S nearest neighbours, towards six of the eight faces of a distorted octahedron. The material is a semiconductor and the Fe ions should be considered to be in a divalent state (as shown by Mössbauer spectroscopy as well as XPS), rather than a tetravalent state as the stoichiometry would suggest. The material as a whole acts as a Van Vleck paramagnet, despite its low-spin divalency.

The positions of X ions in the pyrite structure may be derived from the fluorite structure, starting from a hypothetical Fe ²⁺(S −) ₂ structure. Whereas F− ions in CaF ₂ occupy the centre positions of the eight subcubes of the cubic unit cell (​ ​ ​) etc., the S − ions in FeS ₂ are shifted from these high symmetry positions along <111> axes to reside on ( uuu) and symmetry-equivalent positions. Here, the parameter u should be regarded as a free atomic parameter that takes different values in different pyrite-structure compounds (iron pyrite FeS ₂: u(S) = 0.385). The shift from fluorite u = 0.25 to pyrite u = 0.385 is rather large and creates an S-S distance that is clearly a binding one. This is not surprising as in contrast to F − an ion S − is not a closed shell species. It is isoelectronic with a chlorine atom, also undergoing pairing to form Cl ₂ molecules. Both low spin Fe ²⁺ and the disulfide S ₂ 2− moeties are closed shell entities, explaining the diamagnetic and semiconducting properties.

The S atoms have bonds with three Fe and one other S atom. The site symmetry at Fe and S positions is accounted for by point symmetry groups C 3i and C ₃, respectively. The missing center of inversion at S lattice sites has important consequences for the crystallographic and physical properties of iron pyrite. These consequences derive from the crystal electric field active at the sulfur lattice site, which causes a polarisation of S ions in the pyrite lattice. The polarisation can be calculated on the basis of higher-order Madelung constants and has to be included in the calculation of the lattice energy by using a generalised Born-Haber cycle. This reflects the fact that the covalent bond in the sulfur pair is inadequately accounted for by a strictly ionic treatment.

Arsenopyrite has a related structure with heteroatomic As-S pairs rather than homoatomic ones. Marcasite also possesses homoatomic anion pairs, but the arrangement of the metal and diatomic anions is different from that of pyrite. Despite its name a chalcopyrite () does not contain dianion pairs, but single S 2− sulfide anions.

Pyrite usually forms cuboid crystals, sometimes forming in close association to form raspberry-shaped masses called framboids. However, under certain circumstances, it can form anastamozing filaments or T-shaped crystals. Pyrite can also form shapes almost the same as a regular dodecahedral, known as pyritohedra, and this suggests an explanation for the artificial geometrical models found in Europe as early as the 5th century BC.[clarification needed ]

Cattierite ( Co S ₂) and vaesite ( Ni S ₂) are similar in their structure and belong also to the pyrite group.

Bravoite is a nickel-cobalt bearing variety of pyrite, with > 50% substitution of Ni ²⁺ for Fe ²⁺ within pyrite. Bravoite is not a formally recognised mineral, and is named after the Peruvian scientist Jose J. Bravo (1874-1928).

It is distinguishable from native gold by its hardness, brittleness and crystal form. Natural gold tends to be anhedral (irregularly shaped), whereas pyrite comes as either cubes or multifaceted crystals. Pyrite can often be distinguished by the striations which, in many cases, can be seen on its surface. Chalcopyrite is brighter yellow with a greenish hue when wet and is softer (3.5-4 on Mohs' scale). Arsenopyrite is silver white and does not become more yellow when wet.

Iron pyrite is unstable at Earth's surface: iron pyrite exposed to atmospheric oxygen and water decomposes into iron oxides and sulfate. This process is accelerated by the action of Acidithiobacillus bacteria which oxidize pyrite to produce ferrous iron, sulfate, and protons (). These reactions occur more rapidly when pyrite is finely dispersed (framboidal crystals initially formed by sulfate reducing bacteria (SRB) in argillaceous sediments or dust from mining operations).

Pyrite oxidation and acid mine drainage

Sulfate released from decomposing pyrite combines with water, producing sulfuric acid, leading to acid mine drainage. An example of acid rock drainage caused by pyrite is the 2015 Gold King Mine waste water spill.

.[41]

Dust explosions

Pyrite oxidation is sufficiently exothermic that underground coal mines in high-sulfur coal seams have occasionally had serious problems with spontaneous combustion in the mined-out areas of the mine. The solution is to hermetically seal the mined-out areas to exclude oxygen.[citation needed ]

In modern coal mines, limestone dust is sprayed onto the exposed coal surfaces to reduce the hazard of dust explosions. This has the secondary benefit of neutralizing the acid released by pyrite oxidation and therefore slowing the oxidation cycle described above, thus reducing the likelihood of spontaneous combustion. In the long term, however, oxidation continues, and the hydrated sulfates formed may exert crystallization pressure that can expand cracks in the rock and lead eventually to roof fall.

Weakened building materials

Building stone containing pyrite tends to stain brown as pyrite oxidizes. This problem appears to be significantly worse if any marcasite is present. The presence of pyrite in the aggregate used to make concrete can lead to severe deterioration as pyrite oxidizes. In early 2009, problems with Chinese drywall imported into the United States after Hurricane Katrina were attributed to pyrite oxidation, followed by microbial sulfate reduction which released hydrogen sulfide gas. These problems included a foul odor and corrosion of copper wiring. In the United States, in Canada, and more recently in Ireland, where it was used as underfloor infill, pyrite contamination has caused major structural damage. Normalized tests for aggregate materials certify such materials as free of pyrite.

Pyrite and marcasite commonly occur as replacement pseudomorphs after fossils in black shale and other sedimentary rocks formed under reducing environmental conditions. However, pyrite dollars or pyrite suns which have an appearance similar to sand dollars are pseudofossils and lack the pentagonal symmetry of the animal.

1. ^ Anthony, John W.; Bideaux, Richard A.; Bladh, Kenneth W.; Nichols, Monte C., eds. (1990). "Pyrite" (PDF). Handbook of Mineralogy. Volume I (Elements, Sulfides, Sulfosalts). Chantilly, VA, US: Mineralogical Society of America. ISBN 978-0962209734.

2. ^ Jackson, Julia A.; Mehl, James; Neuendorf, Klaus (2005). Glossary of Geology. American Geological Institute. p. 82. ISBN 9780922152766 - via Google Books.

3. ^ "De re metallica". The Mining Magazine. Translated by Hoover, H.C.; Hoover, L.H. London: Dover. 1950 [1912]. see footnote on p 112.

4. ^ Fleet, M. E.; Mumin, A. Hamid (1997). "Gold-bearing arsenian pyrite and marcasite and arsenopyrite from Carlin Trend gold deposits and laboratory synthesis" (PDF). American Mineralogist. 82 (1-2): 182-193. doi:10.2138/am-1997-1-220. S2CID 55899431.

5. ^ "Cylindrical Primary Lithium [battery]". Lithium-Iron Disulfide (Li-FeS₂) (PDF). Handbook and Application Manual. Energizer Corporation. 2017-09-19.

6. ^ Ellmer, K. & Tributsch, H. (2000-03-11). "Iron Disulfide (Pyrite) as Photovoltaic Material: Problems and Opportunities". Proceedings of the 12th Workshop on Quantum Solar Energy Conversion - (QUANTSOL 2000). Archived from the original on 2010-01-15.

7. Xin Zhang & Mengquin Li (2017-06-19). "Potential resolution to the doping puzzle in iron pyrite: Carrier type determination by Hall effect and thermopower". Physical Review Materials. 1. doi:10.1103/PhysRevMaterials.1.015402.

8. ^ The Principles Underlying Radio Communication. U.S. Army Signal Corps. Radio Pamphlet. 40. 1918. section 179, pp 302-305 - via Google Books.

9. ^ Thomas H. Lee (2004). The Design of Radio Frequency Integrated Circuits (2nd ed.). Cambridge, UK: Cambridge University Press. pp. 4-6. ISBN 9780521835398 - via Google Books.

10. ^ Wadia, Cyrus; Alivisatos, A. Paul; Kammen, Daniel M. (2009). "Materials availability expands the opportunity for large-scale photovoltaics deployment". Environmental Science & Technology. 43 (6): 2072-7. Bibcode:2009EnST...43.2072W. doi:10.1021/es8019534. PMID 19368216. S2CID 36725835.

11. ^ Sanders, Robert (17 February 2009). "Cheaper materials could be key to low-cost solar cells". Berkeley, CA: University of California - Berkeley.

12. ^ Calvo, Miguel and Sevillano, Emilia (1998). "Pyrite crystals from Soria and La Rioja provinces, Spain". The Mineralogical Record. 20: 451-456.

13. ^ Walter, Jeff; Voigt, Bryan; Day-Roberts, Ezra; Heltemes, Kei; Fernandes, Rafael M.; Birol, Turan; Leighton, Chris (1 July 2020). "Voltage-induced ferromagnetism in a diamagnet". Science Advances. 6 (31): eabb7721. doi:10.1126/sciadv.abb7721. ISSN 2375-2548. PMC . PMID 32832693.

14. ^ Kaur, Harneet; Tian, Ruiyuan; Roy, Ahin; McCrystall, Mark; Horváth, Dominik V.; Onrubia, Guillermo L.; Smith, Ross; Ruether, Manuel; Griffin, Aideen; Backes, Claudia; Nicolosi, Valeria; Coleman, Jonathan N. (22 September 2020). "Production of Quasi-2D Platelets of Non-Layered Iron Pyrite (FeS₂) by Liquid-Phase Exfoliation for High Performance Battery Electrodes". ACS Nano. doi:10.1021/acsnano.0c05292.

15. ^ Vaughan, D. J.; Craig, J. R. (1978). Mineral Chemistry of Metal Sulfides. Cambridge, UK: Cambridge University Press. ISBN 978-0-521-21489-6.

16. ^ Bragg, W. L. (1913). "The structure of some crystals as indicated by their diffraction of X-rays". Proceedings of the Royal Society A. 89 (610): 248-277. Bibcode:1913RSPSA..89..248B. doi:10.1098/rspa.1913.0083. JSTOR 93488.

17. ^ Birkholz, M.; Fiechter, S.; Hartmann, A.; Tributsch, H. (1991). "Sulfur deficiency in iron pyrite (FeS2−x) and its consequences for band structure models". Physical Review B. 43 (14): 11926-11936. Bibcode:1991PhRvB..4311926B. doi:10.1103/PhysRevB.43.11926. PMID 9996968.

18. ^ Brese, Nathaniel E.; von Schnering, Hans Georg (1994). "Bonding Trends in Pyrites and a Reinvestigation of the Structure of PdAs₂, PdSb₂, PtSb₂ and PtBi₂". Z. Anorg. Allg. Chem. 620 (3): 393-404. doi:10.1002/zaac.19946200302.

19. ^ Burgardt, P.; Seehra, M. S. (1977-04-01). "Magnetic susceptibility of iron pyrite (FeS2) between 4.2 and 620 K". Solid State Communications. 22 (2): 153-156. doi:10.1016/0038-1098(77)90422-7. ISSN 0038-1098.

20. ^ Stevens, E. D.; Delucia, M. L.; Coppens, P. (1980). "Experimental observation of the Effect of Crystal Field Splitting on the Electron Density Distribution of Iron Pyrite". Inorg. Chem. 19 (4): 813-820. doi:10.1021/ic50206a006.

21. ^ Bonev, I. K.; Garcia-Ruiz, J. M.; Atanassova, R.; Otalora, F.; Petrussenko, S. (2005). "Genesis of filamentary pyrite associated with calcite crystals". European Journal of Mineralogy. 17 (6): 905-913. Bibcode:2005EJMin..17..905B. CiteSeerX 10.1.1.378.3304. doi:10.1127/0935-1221/2005/0017-0905.

22. The pyritohedral form is described as a dodecahedron with ^ pyritohedral symmetry; Dana J. et al., (1944), System of mineralogy, New York, p 282

23. Andrew Roy, Coal Mining in Iowa, ^ Coal Trade Journal, quoted in History of Lucas County Iowa, State Historical Company, Des Moines (1881) pp. 613-615.

24. ^ Zodrow, E (2005). "Colliery and surface hazards through coal-pyrite oxidation (Pennsylvanian Sydney Coalfield, Nova Scotia, Canada)". International Journal of Coal Geology. 64 (1-2): 145-155. doi:10.1016/j.coal.2005.03.013.

25. ^ Tagnithamou, A; Sariccoric, M; Rivard, P (2005). "Internal deterioration of concrete by the oxidation of pyrrhotitic aggregates". Cement and Concrete Research. 35: 99-107. doi:10.1016/j.cemconres.2004.06.030.

26. "^ PYRITE and Your House, What Home-Owners Should Know Archived 2012-01-06 at the Wayback Machine" - ISBN 2-922677-01-X - Legal deposit - National Library of Canada, May 2000

27. Shrimer, F. and Bromley, AV (2012) "Pyritic Heave in Ireland". ^ Proceedings of the Euroseminar on Building Materials. International Cement Microscopy Association (Halle Germany)

28. Brennan, Michael (22 February 2010) ^ Devastating 'pyrite epidemic' hits 20,000 newly built houses. Irish Independent

29. ^ Briggs, D. E. G.; Raiswell, R.; Bottrell, S. H.; Hatfield, D.; Bartels, C. (1996-06-01). "Controls on the pyritization of exceptionally preserved fossils; an analysis of the Lower Devonian Hunsrueck Slate of Germany". American Journal of Science. 296 (6): 633-663. Bibcode:1996AmJS..296..633B. doi:10.2475/ajs.296.6.633. ISSN 0002-9599.

American Geological Institute, 2003, Dictionary of Mining, Mineral, and Related Terms, 2nd ed., Springer, New York, ISBN 978-3-540-01271-9.

David Rickard, Pyrite: A Natural History of Fool's Gold, Oxford, New York, 2015, ISBN 978-0-19-020367-2.

# data[].attributes.attachment_link.description
The mineral pyrite (), or iron pyrite, also known as fool's gold, is an iron sulfide with the chemical formula Fe S 2 (iron (II) disulfide). Pyrite is the most abundant sulfide mineral. Pyrite's metallic luster and pale brass-yellow hue give it a superficial resemblance to gold, hence the well-known nickname of fool's gold.

# data[].attributes.body
Its luster is non metallic
Has a colorless appearance
Produces a white streak
Cleavage
It takes the strength of a glass plate to create a scratch
Used to make cement

# data[].attributes.attachment_link.description
The uses and properties of the mineral calcite with numerous photos.