Chalcogen: Difference between revisions
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The '''chalcogens''' | The '''chalcogens''' ({{IPAc-en|ˈ|k|æ|l|k|ə|dʒ|ə|n|z}}, {{Respell|KAL|kə|jənz}}) are the [[chemical element]]s in [[group (periodic table)|group]] 16 of the [[periodic table]].<ref>{{Cite book |last1=House |first1=James E. |title=Inorganic chemistry |last2=House |first2=James Evan |date=2008 |publisher=Elsevier Academic Press |isbn=978-0-12-356786-4 |location=Amsterdam Heidelberg |page=523}}</ref> This group is also known as the '''oxygen family'''. Group 16 consists of the elements [[oxygen]] (O), [[sulfur]] (S), [[selenium]] (Se), [[tellurium]] (Te), and the [[Radioactive decay|radioactive]] elements [[polonium]] (Po) and [[livermorium]] (Lv).<ref name="ReferenceB">{{cite book|last=Emsley|first=John|title=Nature's Building Blocks: An A-Z Guide to the Elements|edition=New|year=2011|publisher=Oxford University Press|location=New York, NY|isbn=978-0-19-960563-7|pages=375–383, 412–415, 475–481, 511–520, 529–533, 582}}</ref> Often, oxygen is treated separately from the other chalcogens, sometimes even excluded from the scope of the term "chalcogen" altogether, due to its very different chemical behavior from sulfur, selenium, tellurium, and polonium. The word "chalcogen" means "[[ore]]-forming"; chalcogens got their name because [[protoscience|protoscientists]] and early scientists could discern that these essences (which science would later reveal to be chemical elements) were involved in ore formation.<!--You are in the lede of the article. Do not add the arcana of the stymology here. It is already covered in the [[#Names and etymology]] section. The purpose of this sentence being here is summarization of the point for the lede's purposes, and it is already optimized for that.--> | ||
Sulfur has been known since antiquity, and oxygen was recognized as an element in the 18th century. Selenium, tellurium and polonium were discovered in the 19th century, and livermorium in 2000. All of the chalcogens have six [[valence electron]]s, leaving them two electrons short of a full outer shell. Their most common [[oxidation state]]s are −2, +2, +4, and +6. They have relatively small [[atomic radius|atomic radii]], especially the lighter ones.<ref name="Jackson2002" /> | Sulfur has been known since antiquity, and oxygen was recognized as an element in the 18th century. Selenium, tellurium and polonium were discovered in the 19th century, and livermorium in 2000. All of the chalcogens have six [[valence electron]]s, leaving them two electrons short of a full outer shell. Their most common [[oxidation state]]s are −2, +2, +4, and +6. They have relatively small [[atomic radius|atomic radii]], especially the lighter ones.<ref name="Jackson2002" /> | ||
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Sulfur has over 20 known allotropes, which is more than any other element except [[allotropes of carbon|carbon]].<ref name="Greenwood">{{Greenwood&Earnshaw|pages = 645–662}}</ref> The most common allotropes are in the form of eight-atom rings, but other molecular allotropes that contain as few as two atoms or as many as 20 are known. Other notable sulfur allotropes include [[rhombic crystal system|rhombic]] sulfur and [[monoclinic]] sulfur. Rhombic sulfur is the more stable of the two allotropes. Monoclinic sulfur takes the form of long needles and is formed when liquid sulfur is cooled to slightly below its melting point. The atoms in liquid sulfur are generally in the form of long chains, but above 190 °C, the chains begin to break down. If liquid sulfur above 190 °C is [[freezing|frozen]] very rapidly, the resulting sulfur is amorphous or "plastic" sulfur. Gaseous sulfur is a mixture of diatomic sulfur (S<sub>2</sub>) and 8-atom rings.<ref>{{cite web|last = McClure|first = Mark R.|url = http://www.uncp.edu/home/mcclurem/ptable/sulfur/s.htm|title = sulfur|access-date = November 25, 2013|archive-url = https://web.archive.org/web/20140312220122/http://www2.uncp.edu/home/mcclurem/ptable/sulfur/s.htm|archive-date = March 12, 2014|url-status = dead|df = mdy-all}}</ref> | Sulfur has over 20 known allotropes, which is more than any other element except [[allotropes of carbon|carbon]].<ref name="Greenwood">{{Greenwood&Earnshaw|pages = 645–662}}</ref> The most common allotropes are in the form of eight-atom rings, but other molecular allotropes that contain as few as two atoms or as many as 20 are known. Other notable sulfur allotropes include [[rhombic crystal system|rhombic]] sulfur and [[monoclinic]] sulfur. Rhombic sulfur is the more stable of the two allotropes. Monoclinic sulfur takes the form of long needles and is formed when liquid sulfur is cooled to slightly below its melting point. The atoms in liquid sulfur are generally in the form of long chains, but above 190 °C, the chains begin to break down. If liquid sulfur above 190 °C is [[freezing|frozen]] very rapidly, the resulting sulfur is amorphous or "plastic" sulfur. Gaseous sulfur is a mixture of diatomic sulfur (S<sub>2</sub>) and 8-atom rings.<ref>{{cite web|last = McClure|first = Mark R.|url = http://www.uncp.edu/home/mcclurem/ptable/sulfur/s.htm|title = sulfur|access-date = November 25, 2013|archive-url = https://web.archive.org/web/20140312220122/http://www2.uncp.edu/home/mcclurem/ptable/sulfur/s.htm|archive-date = March 12, 2014|url-status = dead|df = mdy-all}}</ref> | ||
Selenium has at least eight distinct allotropes.<ref>{{Greenwood&Earnshaw2nd|page=751}}</ref> The gray allotrope, commonly referred to as the "metallic" allotrope, despite not being a metal, is stable and has a hexagonal [[crystal structure]]. The gray allotrope of selenium is soft, with a [[Mohs hardness]] of 2, and brittle. Four other allotropes of selenium are [[metastable]]. These include two [[monoclinic]] red allotropes and two [[amorphous]] allotropes, one of which is red and one of which is black.<ref>{{Cite web|vauthors = Butterman WC, ((Brown RD Jr)) |url = | Selenium has at least eight distinct allotropes.<ref>{{Greenwood&Earnshaw2nd|page=751}}</ref> The gray allotrope, commonly referred to as the "metallic" allotrope, despite not being a metal, is stable and has a hexagonal [[crystal structure]]. The gray allotrope of selenium is soft, with a [[Mohs hardness]] of 2, and brittle. Four other allotropes of selenium are [[metastable]]. These include two [[monoclinic]] red allotropes and two [[amorphous]] allotropes, one of which is red and one of which is black.<ref>{{Cite web|vauthors = Butterman WC, ((Brown RD Jr)) |url =https://pubs.usgs.gov/of/2003/of03-018/of03-018.pdf |title = Selenium. Mineral Commodity Profiles|year = 2004 |publisher = Department of the Interior |url-status=live |archive-url=https://web.archive.org/web/20121003211018/http://pubs.usgs.gov/of/2003/of03-018/of03-018.pdf |archive-date=October 3, 2012 |access-date=November 25, 2013}}</ref> The red allotrope converts to the black allotrope in the presence of heat. The gray allotrope of selenium is made from [[spiral]]s on selenium atoms, while one of the red allotropes is made of stacks of selenium rings (Se<sub>8</sub>).<ref name="ReferenceB"/>{{dubious|date=September 2014}} | ||
Tellurium is not known to have any allotropes,<ref>{{cite web |last = Emsley |first = John |url = http://www.rsc.org/periodic-table/element/52/tellurium |title = Tellurium |year = 2011 |access-date=November 25, 2013 |publisher=Royal Society of Chemistry}}</ref> although its typical form is hexagonal. Polonium has two allotropes, which are known as α-polonium and β-polonium.<ref>{{cite web|last = Emsley|first = John|url =http://www.rsc.org/periodic-table/element/84/polonium|title = Polonium|year = 2011 |access-date=November 25, 2013 |publisher=Royal Society of Chemistry}}</ref> α-polonium has a cubic crystal structure and converts to the rhombohedral β-polonium at 36 °C.<ref name="ReferenceB"/> | Tellurium is not known to have any allotropes,<ref>{{cite web |last = Emsley |first = John |url = http://www.rsc.org/periodic-table/element/52/tellurium |title = Tellurium |year = 2011 |access-date=November 25, 2013 |publisher=Royal Society of Chemistry}}</ref> although its typical form is hexagonal. Polonium has two allotropes, which are known as α-polonium and β-polonium.<ref>{{cite web|last = Emsley|first = John|url =http://www.rsc.org/periodic-table/element/84/polonium|title = Polonium|year = 2011 |access-date=November 25, 2013 |publisher=Royal Society of Chemistry}}</ref> α-polonium has a cubic crystal structure and converts to the rhombohedral β-polonium at 36 °C.<ref name="ReferenceB"/> | ||
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Sulfur's oxidation states are −2, +2, +4, and +6. Sulfur-containing analogs of oxygen compounds often have the prefix ''thio-''. Sulfur's chemistry is similar to oxygen's, in many ways. One difference is that sulfur-sulfur [[double bond]]s are far weaker than oxygen-oxygen double bonds, but sulfur-sulfur [[single bond]]s are stronger than oxygen-oxygen single bonds.<ref>{{cite web|url = http://chemed.chem.purdue.edu/genchem/topicreview/bp/ch10/group6.php#selenium|title = The Chemistry of Oxygen and Sulfur |access-date=November 25, 2013 |publisher = Bodner Research Web}}</ref> Organic sulfur compounds such as [[thiol]]s have a strong specific smell, and a few are utilized by some organisms.<ref name="ReferenceB"/> | Sulfur's oxidation states are −2, +2, +4, and +6. Sulfur-containing analogs of oxygen compounds often have the prefix ''thio-''. Sulfur's chemistry is similar to oxygen's, in many ways. One difference is that sulfur-sulfur [[double bond]]s are far weaker than oxygen-oxygen double bonds, but sulfur-sulfur [[single bond]]s are stronger than oxygen-oxygen single bonds.<ref>{{cite web|url = http://chemed.chem.purdue.edu/genchem/topicreview/bp/ch10/group6.php#selenium|title = The Chemistry of Oxygen and Sulfur |access-date=November 25, 2013 |publisher = Bodner Research Web}}</ref> Organic sulfur compounds such as [[thiol]]s have a strong specific smell, and a few are utilized by some organisms.<ref name="ReferenceB"/> | ||
Selenium's oxidation states are −2, +4, and +6. Selenium, like most chalcogens, bonds with oxygen.<ref name="ReferenceB"/> There are some [[organoselenium chemistry|organic selenium compounds]], such as [[ | Selenium's oxidation states are −2, +4, and +6. Selenium, like most chalcogens, bonds with oxygen.<ref name="ReferenceB"/> There are some [[organoselenium chemistry|organic selenium compounds]], such as [[selenoprotein]]s. Tellurium's oxidation states are −2, +2, +4, and +6.<ref name="Jackson2002" /> Tellurium forms the oxides [[tellurium monoxide]], [[tellurium dioxide]], and [[tellurium trioxide]].<ref name="ReferenceB"/> Polonium's oxidation states are +2 and +4.<ref name="Jackson2002" /> | ||
[[File:Brindis (24675281395).jpg|thumb|upright|left|[[Water]] ({{chem2|H2O}}) is the most familiar chalcogen-containing compound.|alt=Water dripping into a glass, showing drops and bubbles.]] | [[File:Brindis (24675281395).jpg|thumb|upright|left|[[Water]] ({{chem2|H2O}}) is the most familiar chalcogen-containing compound.|alt=Water dripping into a glass, showing drops and bubbles.]] | ||
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===With halogens=== <!-- Chalcohalide redirects here --> | ===With halogens=== <!-- Chalcohalide redirects here --> | ||
Chalcogens also form compounds with [[halogen]]s known as '''chalcohalides''', or '''chalcogen halides'''. The majority of simple chalcogen halides are well-known and widely used as chemical [[reagent]]s. However, more complicated chalcogen halides, such as sulfenyl, sulfonyl, and sulfuryl halides, are less well known to science. Out of the compounds consisting purely of chalcogens and halogens, there are a total of 13 chalcogen fluorides, nine chalcogen chlorides, eight chalcogen bromides, and six chalcogen iodides that are known.{{dubious|date=September 2014}} The heavier chalcogen halides often have significant molecular interactions. Sulfur fluorides with low valences are fairly unstable and little is known about their properties.{{dubious|date=September 2014}} However, sulfur fluorides with high valences, such as [[sulfur hexafluoride]], are stable and well-known. [[Sulfur tetrafluoride]] is also a well-known sulfur fluoride. Certain selenium fluorides, such as [[selenium difluoride]], have been produced in small amounts. The crystal structures of both [[selenium tetrafluoride]] and [[tellurium tetrafluoride]] are known. Chalcogen chlorides and bromides have also been explored. In particular, selenium dichloride and sulfur dichloride can react to form [[Organoselenium chemistry|organic selenium compounds]]. Dichalcogen dihalides, such as Se<sub>2</sub>Cl<sub>2</sub> also are known to exist. There are also mixed chalcogen-halogen compounds. These include SeSX, with X being chlorine or bromine.{{ | Chalcogens also form compounds with [[halogen]]s known as '''chalcohalides''', or '''chalcogen halides'''. The majority of simple chalcogen halides are well-known and widely used as chemical [[reagent]]s. However, more complicated chalcogen halides, such as sulfenyl, sulfonyl, and sulfuryl halides, are less well known to science. Out of the compounds consisting purely of chalcogens and halogens, there are a total of 13 chalcogen fluorides, nine chalcogen chlorides, eight chalcogen bromides, and six chalcogen iodides that are known.{{dubious|date=September 2014}} The heavier chalcogen halides often have significant molecular interactions. Sulfur fluorides with low valences are fairly unstable and little is known about their properties.{{dubious|date=September 2014}} However, sulfur fluorides with high valences, such as [[sulfur hexafluoride]], are stable and well-known. [[Sulfur tetrafluoride]] is also a well-known sulfur fluoride. Certain selenium fluorides, such as [[selenium difluoride]], have been produced in small amounts. The crystal structures of both [[selenium tetrafluoride]] and [[tellurium tetrafluoride]] are known. Chalcogen chlorides and bromides have also been explored. In particular, selenium dichloride and sulfur dichloride can react to form [[Organoselenium chemistry|organic selenium compounds]]. Dichalcogen dihalides, such as Se<sub>2</sub>Cl<sub>2</sub> also are known to exist. There are also mixed chalcogen-halogen compounds. These include SeSX<sub>2</sub>, with X being chlorine or bromine.<ref name="Milne">{{cite journal |last1=Milne |first1=John B. |title=Selenium sulfur dihalides, ChnX2 (n = 1, 2, 3; Ch = Se, S; X = Br, Cl). Raman and 77Se NMR spectroscopic characterization |journal=Canadian Journal of Chemistry |date=March 1992 |volume=70 |issue=3 |pages=693–699 |doi=10.1139/v92-092 |url=https://cdnsciencepub.com/doi/10.1139/v92-092 |issn=0008-4042|url-access=subscription }}</ref> Such compounds can form in mixtures of [[sulfur dichloride]] and selenium halides.<ref name="Milne"/> These compounds have been fairly recently structurally characterized, as of 2008. In general, diselenium and disulfur chlorides and bromides are useful chemical reagents. Chalcogen halides with attached metal atoms are soluble in organic solutions.{{dubious|date=September 2014}} One example of such a compound is {{chem2|[[Mo]]S2Cl3}}. Unlike selenium chlorides and bromides, selenium [[iodide]]s have not been isolated, as of 2008, although it is likely that they occur in solution. Diselenium diiodide, however, does occur in equilibrium with selenium atoms and iodine molecules. Some tellurium halides with low valences, such as {{chem2|Te2Cl2}} and {{chem2|Te2Br2}}, form [[polymer]]s when in the [[Solid-state chemistry|solid state]]. These tellurium halides can be synthesized by the reduction of pure tellurium with [[superhydride]] and reacting the resulting product with tellurium tetrahalides. Ditellurium dihalides tend to get less stable as the halides become lower in atomic number and atomic mass. Tellurium also forms iodides with even fewer iodine atoms than diiodides. These include TeI and Te<sub>2</sub>I. These compounds have extended structures in the solid state. Halogens and chalcogens can also form halochalcogenate [[anion]]s.<ref name="handbook"/> | ||
===Organic=== | ===Organic=== | ||
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===Names and etymology=== | ===Names and etymology=== | ||
In the 19th century, [[Jons Jacob Berzelius]] suggested calling the elements in group 16 "amphigens",<ref name="che.uc">{{cite journal|url=http://www.che.uc.edu/jensen/W.%20B.%20Jensen/Reprints/072.%20Chalcogen.pdf|title=A Note on the Term "Chalcogen"|doi=10.1021/ed074p1063|year=1997|author1-link=William B. Jensen|last1=Jensen|first1=William B.|journal=Journal of Chemical Education|volume=74|issue=9|pages=1063|bibcode=1997JChEd..74.1063J|access-date=November 25, 2013|archive-date=October 29, 2013|archive-url=https://web.archive.org/web/20131029185247/http://www.che.uc.edu/jensen/W.%20B.%20Jensen/Reprints/072.%20Chalcogen.pdf|url-status=dead}}</ref> as the elements in the group formed [[amphid salts]] (salts of [[oxyacid]]s,<ref>{{cite web |url=http://dictionary.reference.com/browse/oxysalt |title=Oxysalt - Define Oxysalt at Dictionary.com |publisher=Dictionary.reference.com |access-date=November 25, 2013}}</ref><ref>{{cite web |url=http://www.thefreedictionary.com/Amphigen |title=Amphigen – definition of Amphigen by the Free Online Dictionary, Thesaurus and Encyclopedia |publisher=Thefreedictionary.com |access-date=November 25, 2013}}</ref> formerly regarded as composed of two oxides, an acid and a basic oxide). The term received some use in the early 1800s but is now obsolete.<ref name="che.uc"/> The name ''chalcogen'' comes from the Greek words ''{{lang|grc|χαλκος}}'' ({{lang|grc-Latn|chalkos}}, | In the 19th century, [[Jons Jacob Berzelius]] suggested calling the elements in group 16 "amphigens",<ref name="che.uc">{{cite journal|url=http://www.che.uc.edu/jensen/W.%20B.%20Jensen/Reprints/072.%20Chalcogen.pdf|title=A Note on the Term "Chalcogen"|doi=10.1021/ed074p1063|year=1997|author1-link=William B. Jensen|last1=Jensen|first1=William B.|journal=Journal of Chemical Education|volume=74|issue=9|pages=1063|bibcode=1997JChEd..74.1063J|access-date=November 25, 2013|archive-date=October 29, 2013|archive-url=https://web.archive.org/web/20131029185247/http://www.che.uc.edu/jensen/W.%20B.%20Jensen/Reprints/072.%20Chalcogen.pdf|url-status=dead}}</ref> as the elements in the group formed [[amphid salts]] (salts of [[oxyacid]]s,<ref>{{cite web |url=http://dictionary.reference.com/browse/oxysalt |title=Oxysalt - Define Oxysalt at Dictionary.com |publisher=Dictionary.reference.com |access-date=November 25, 2013}}</ref><ref>{{cite web |url=http://www.thefreedictionary.com/Amphigen |title=Amphigen – definition of Amphigen by the Free Online Dictionary, Thesaurus and Encyclopedia |publisher=Thefreedictionary.com |access-date=November 25, 2013}}</ref> formerly regarded as composed of two oxides, an acid and a basic oxide). The term received some use in the early 1800s but is now obsolete.<ref name="che.uc"/> The name ''chalcogen'' comes from the Greek words ''{{lang|grc|χαλκος}}'' ({{lang|grc-Latn|chalkos}}, most narrowly "[[copper]]", but the term was also used for [[bronze]], [[brass]], any metal in the poetic sense, [[ore]], and [[coin]]<ref name="Chalco-">{{Cite book|title=The New Shorter Oxford Dictionary|year=1993|publisher=Oxford University Press|isbn=978-0-19-861134-9|page=[https://archive.org/details/newshorteroxford00lesl/page/368 368]|url=https://archive.org/details/newshorteroxford00lesl/page/368}}</ref>), and ''{{lang|grc|γενές}}'' ({{lang|grc-Latn|genes}}, born,<ref>{{cite web |last=Harper|first=Douglas|title=Online Etymology Dictionary|url=http://www.etymonline.com/index.php?term=-gen |access-date=November 25, 2013}}</ref> gender, kindle, produce).<ref>{{cite encyclopedia |url = http://www.merriam-webster.com/dictionary/chalcogen |title= chalcogen|year = 2013 |access-date=November 25, 2013 |dictionary=Merriam-Webster}}</ref><ref>{{Cite book|author = Bouroushian, M.|url = https://books.google.com/books?id=B8WgWHjN54oC&pg=PA1|title = Electrochemistry of Metal Chalcogenides|year = 2010|isbn = 978-3-642-03967-6|bibcode = 2010emc..book.....B|doi = 10.1007/978-3-642-03967-6|series = Monographs in Electrochemistry }}</ref> It was first used in 1932 by [[Wilhelm Biltz]]'s group at [[Leibniz University Hannover]], where it was proposed by {{ill|Werner Fischer (chemist)|lt=Werner Fischer|de|Werner Fischer (Chemiker)}}.<ref name="chalcogen2">{{cite journal|author=Fischer, Werner|title=A Second Note on the Term "Chalcogen"|journal=Journal of Chemical Education|year=2001|volume=78|issue=10|page=1333|doi=10.1021/ed078p1333.1|bibcode = 2001JChEd..78.1333F |doi-access=}}</ref> The word "chalcogen" gained popularity in Germany during the 1930s because the term was analogous to "halogen".<ref>{{cite book|author=Krebs, Robert E. |title=The History And Use of Our Earth's Chemical Elements: A Reference Guide |url=https://books.google.com/books?id=yb9xTj72vNAC&pg=PA223 |year=2006 |publisher=Greenwood Publishing Group |isbn=978-0-313-33438-2 |pages=223– |access-date=November 25, 2013}}</ref> Although the literal meanings of the modern Greek words imply that ''chalcogen'' means "copper-former", this is misleading because the chalcogens have nothing to do with copper in particular. "Ore-former" has been suggested as a better translation,<ref name="chalcogen">{{cite journal|author=Jensen, William B.|journal=Journal of Chemical Education|year=1997|volume=74|issue=9|page=1063|doi=10.1021/ed074p1063|title=A Note on the Term "Chalcogen"|bibcode = 1997JChEd..74.1063J}}</ref> as the vast majority of metal ores are chalcogenides and the word ''{{lang|grc|χαλκος}}'' in ancient Greek was associated with metals and metal-bearing rock in general; copper, and its alloy [[bronze]], was one of the first metals to be used by humans. | ||
Oxygen's name comes from the Greek words ''oxy genes'', meaning "acid-forming". Sulfur's name comes from either the Latin word ''{{lang|la|sulfurium}}'' or the [[Sanskrit]] word ''{{lang|sa-Latn|sulvere}}''; both of those terms are ancient words for sulfur. Selenium is named after the Greek goddess of the moon, [[Selene]], to match the previously discovered element tellurium, whose name comes from the Latin word ''{{lang|la|telus}}'', meaning earth. Polonium is named after Marie Curie's country of birth, Poland.<ref name="The Elements"/> Livermorium is named for the [[Lawrence Livermore National Laboratory]].<ref>{{cite web|last = Stark|first = Anne M|url = https://www.llnl.gov/news/newsreleases/2012/May/NR-12-05-07.html|title = Livermorium and Flerovium join the periodic table of elements|date = May 2012|access-date = November 25, 2013|archive-date = February 19, 2013|archive-url = https://web.archive.org/web/20130219040850/https://www.llnl.gov/news/newsreleases/2012/May/NR-12-05-07.html|url-status = dead}}</ref> | Oxygen's name comes from the Greek words ''oxy genes'', meaning "acid-forming". Sulfur's name comes from either the Latin word ''{{lang|la|sulfurium}}'' or the [[Sanskrit]] word ''{{lang|sa-Latn|sulvere}}''; both of those terms are ancient words for sulfur. Selenium is named after the Greek goddess of the moon, [[Selene]], to match the previously discovered element tellurium, whose name comes from the Latin word ''{{lang|la|telus}}'', meaning earth. Polonium is named after Marie Curie's country of birth, Poland.<ref name="The Elements"/> Livermorium is named for the [[Lawrence Livermore National Laboratory]].<ref>{{cite web|last = Stark|first = Anne M|url = https://www.llnl.gov/news/newsreleases/2012/May/NR-12-05-07.html|title = Livermorium and Flerovium join the periodic table of elements|date = May 2012|access-date = November 25, 2013|archive-date = February 19, 2013|archive-url = https://web.archive.org/web/20130219040850/https://www.llnl.gov/news/newsreleases/2012/May/NR-12-05-07.html|url-status = dead}}</ref> | ||
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The world production of selenium is around 1500 metric tons per year, out of which roughly 10% is recycled. Japan is the largest producer, producing 800 metric tons of selenium per year. Other large producers include Belgium (300 metric tons per year), the United States (over 200 metric tons per year), Sweden (130 metric tons per year), and Russia (100 metric tons per year). Selenium can be extracted from the waste from the process of electrolytically refining copper. Another method of producing selenium is to farm selenium-gathering plants such as [[milk vetch]]. This method could produce three kilograms of selenium per acre, but is not commonly practiced.<ref name = "ReferenceB"/> | The world production of selenium is around 1500 metric tons per year, out of which roughly 10% is recycled. Japan is the largest producer, producing 800 metric tons of selenium per year. Other large producers include Belgium (300 metric tons per year), the United States (over 200 metric tons per year), Sweden (130 metric tons per year), and Russia (100 metric tons per year). Selenium can be extracted from the waste from the process of electrolytically refining copper. Another method of producing selenium is to farm selenium-gathering plants such as [[milk vetch]]. This method could produce three kilograms of selenium per acre, but is not commonly practiced.<ref name = "ReferenceB"/> | ||
Tellurium is mostly produced as a by-product of the processing of copper.<ref>{{cite web |author = Callaghan, R. |url = | Tellurium is mostly produced as a by-product of the processing of copper.<ref>{{cite web |author = Callaghan, R. |url = https://minerals.usgs.gov/minerals/pubs/commodity/selenium/ |title = Selenium and Tellurium Statistics and Information |year = 2011 |access-date = November 25, 2013 |publisher = United States Geological Survey |archive-date = May 8, 2012 |archive-url = https://web.archive.org/web/20120508085217/http://minerals.usgs.gov/minerals/pubs/commodity/selenium/ |url-status = dead }}</ref> Tellurium can also be refined by [[electrolytic reduction]] of [[sodium telluride]]. The world production of tellurium is between 150 and 200 metric tons per year. The United States is one of the largest producers of tellurium, producing around 50 metric tons per year. Peru, Japan, and Canada are also large producers of tellurium.<ref name="ReferenceB"/> | ||
Until the creation of nuclear reactors, all polonium had to be extracted from uranium ore. In modern times, most [[isotopes of polonium]] are produced by bombarding [[bismuth]] with neutrons.<ref name="The Elements"/> Polonium can also be produced by high [[neutron flux]]es in [[nuclear reactors]]. Approximately 100 grams of polonium are produced yearly.<ref name = "Factsheets"/> All the polonium produced for commercial purposes is made in the Ozersk nuclear reactor in Russia. From there, it is taken to [[Samara, Russia]] for purification, and from there to [[St. Petersburg]] for distribution. The United States is the largest consumer of polonium.<ref name = "ReferenceB"/> | Until the creation of nuclear reactors, all polonium had to be extracted from uranium ore. In modern times, most [[isotopes of polonium]] are produced by bombarding [[bismuth]] with neutrons.<ref name="The Elements"/> Polonium can also be produced by high [[neutron flux]]es in [[nuclear reactors]]. Approximately 100 grams of polonium are produced yearly.<ref name = "Factsheets"/> All the polonium produced for commercial purposes is made in the Ozersk nuclear reactor in Russia. From there, it is taken to [[Samara, Russia]] for purification, and from there to [[St. Petersburg]] for distribution. The United States is the largest consumer of polonium.<ref name = "ReferenceB"/> | ||
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Oxygen is generally nontoxic, but [[oxygen toxicity]] has been reported when it is used in high concentrations. In both elemental gaseous form and as a component of water, it is vital to almost all life on Earth. Despite this, liquid oxygen is highly dangerous.<ref name = "The Elements"/> Even gaseous oxygen is dangerous in excess. For instance, [[Underwater diving|sports divers]] have occasionally drowned from [[convulsion]]s caused by breathing pure oxygen at a depth of more than {{convert|10|m|ft|abbr=off|sp=us}} underwater.<ref name = "ReferenceB"/> Oxygen is also toxic to some [[bacteria]].<ref name = "Structure of Matter"/> Ozone, an allotrope of oxygen, is toxic to most life. It can cause [[lesion]]s in the respiratory tract.<ref>{{Cite journal|last = Menzel|first = D.B.|pmid=6376815|title = Ozone: an overview of its toxicity in man and animals|year = 1984|doi=10.1080/15287398409530493|journal = Journal of Toxicology and Environmental Health|volume = 13|issue = 2–3|pages = 183–204| bibcode=1984JTEH...13..181M }}</ref> | Oxygen is generally nontoxic, but [[oxygen toxicity]] has been reported when it is used in high concentrations. In both elemental gaseous form and as a component of water, it is vital to almost all life on Earth. Despite this, liquid oxygen is highly dangerous.<ref name = "The Elements"/> Even gaseous oxygen is dangerous in excess. For instance, [[Underwater diving|sports divers]] have occasionally drowned from [[convulsion]]s caused by breathing pure oxygen at a depth of more than {{convert|10|m|ft|abbr=off|sp=us}} underwater.<ref name = "ReferenceB"/> Oxygen is also toxic to some [[bacteria]].<ref name = "Structure of Matter"/> Ozone, an allotrope of oxygen, is toxic to most life. It can cause [[lesion]]s in the respiratory tract.<ref>{{Cite journal|last = Menzel|first = D.B.|pmid=6376815|title = Ozone: an overview of its toxicity in man and animals|year = 1984|doi=10.1080/15287398409530493|journal = Journal of Toxicology and Environmental Health|volume = 13|issue = 2–3|pages = 183–204| bibcode=1984JTEH...13..181M }}</ref> | ||
Sulfur is generally nontoxic and is even a vital nutrient for humans. However, in its elemental form it can cause redness in the eyes and skin, a burning sensation and a cough if inhaled, a burning sensation and diarrhoea and/or [[catharsis (medicine)|catharsis]]<ref name = EPA-sulfur /> if ingested, and can irritate the mucous membranes.<ref>{{Cite web |url=http://npic.orst.edu/factsheets/sulfurgen.html |title=Sulfur General Fact Sheet |website=npic.orst.edu |access-date=2019-01-23}}</ref><ref>{{cite web|author = Extension Toxicology Network |url = | Sulfur is generally nontoxic and is even a vital nutrient for humans. However, in its elemental form it can cause redness in the eyes and skin, a burning sensation and a cough if inhaled, a burning sensation and diarrhoea and/or [[catharsis (medicine)|catharsis]]<ref name = EPA-sulfur /> if ingested, and can irritate the mucous membranes.<ref>{{Cite web |url=http://npic.orst.edu/factsheets/sulfurgen.html |title=Sulfur General Fact Sheet |website=npic.orst.edu |access-date=2019-01-23}}</ref><ref>{{cite web|author = Extension Toxicology Network |url = https://pmep.cce.cornell.edu/profiles/extoxnet/pyrethrins-ziram/sulfur-ext.html |title = Sulfur|date = September 1995 |access-date=November 25, 2013}}</ref> An excess of sulfur can be toxic for [[Cattle|cows]] because microbes in the [[rumen]]s of cows produce toxic hydrogen sulfide upon reaction with sulfur.<ref>{{cite web|author = College of Veterinary Medicine, Iowa State University|url = http://vetmed.iastate.edu/diagnostic-lab/diagnostic-services/diagnostic-sections/chemistry-/-toxicology/polio-cattle-can-be-ca|title = Sulfur Toxicity|year = 2013 |access-date=November 25, 2013}}</ref> Many sulfur compounds, such as [[hydrogen sulfide]] (H<sub>2</sub>S) and [[sulfur dioxide]] (SO<sub>2</sub>) are highly toxic.<ref name="ReferenceB"/> | ||
Selenium is a trace nutrient required by humans on the order of tens or hundreds of micrograms per day. A dose of over 450 micrograms can be toxic, resulting in bad breath and [[body odor]]. Extended, low-level exposure, which can occur at some industries, results in [[weight loss]], [[anemia]], and [[dermatitis]]. In many cases of selenium poisoning, [[selenous acid]] is formed in the body.<ref>{{cite journal|last = Nutall|first = Kern L.|url = http://www.annclinlabsci.org/content/36/4/409.full|title = Evaluating Selenium Poisoning|year = 2006|pmid = 17127727|volume = 36|issue = 4|pages = 409–20|journal = Annals of Clinical and Laboratory Science}}</ref> [[Hydrogen selenide]] (H<sub>2</sub>Se) is highly toxic.<ref name="ReferenceB"/> | Selenium is a trace nutrient required by humans on the order of tens or hundreds of micrograms per day. A dose of over 450 micrograms can be toxic, resulting in bad breath and [[body odor]]. Extended, low-level exposure, which can occur at some industries, results in [[weight loss]], [[anemia]], and [[dermatitis]]. In many cases of selenium poisoning, [[selenous acid]] is formed in the body.<ref>{{cite journal|last = Nutall|first = Kern L.|url = http://www.annclinlabsci.org/content/36/4/409.full|title = Evaluating Selenium Poisoning|year = 2006|pmid = 17127727|volume = 36|issue = 4|pages = 409–20|journal = Annals of Clinical and Laboratory Science}}</ref> [[Hydrogen selenide]] (H<sub>2</sub>Se) is highly toxic.<ref name="ReferenceB"/> | ||