View source for Chemistry:Chalcogen

{{short description|Group of chemical elements}}
{{Infobox periodic table group/header
| title       = Chalcogens
| group number= 16
| trivial name= chalcogens
| by element  = oxygen group
| CAS         = VIA
| old IUPAC   = VIB
| mark        = O,S,Se,Te,Po,Lv
| left        = [[Chemistry:Pnictogen|pnictogens]]
| right       = [[Chemistry:Halogen|halogens]]}}
|-
! colspan=2 style="text-align:left;" | ↓&nbsp;[[Chemistry:Period (periodic table)|Period]]
|-
| [[Chemistry:Period 2 element|'''2''']]
| {{element cell image|8|Oxygen|O| |Gas|Other nonmetal|Primordial|image=Oxygen_discharge_tube.jpg|imaged caption=An oxygen discharge tube}}
|-
! [[Chemistry:Period 3 element|3]]
| {{element cell image|16|Sulfur|S| |Solid|Other nonmetal|Primordial|image=Sulfur-sample.jpg|imge caption=A sample of sulfur}}
|-
! [[Chemistry:Period 4 element|4]]
| {{element cell image|34|Selenium|Se| |Solid|Other nonmetal|Primordial|image=SeBlackRed.jpg|image caption=2 allotropes of selenium: black and red. 3 others not shown.}}
|-
! [[Chemistry:Period 5 element|5]]
| {{element cell image|52|Tellurium|Te| |Solid|Metalloid|Primordial|image=Tellurium2.jpg|image caption=Tellurium in metallic form}}
|-
! [[Chemistry:Period 6 element|6]]
| {{element cell image|84|Polonium|Po| |Solid|Other metal|From decay|image=<!-- Do NOT use a non-free image here, per WP:NFCC #9 -->|image link=File:Polonium.jpg|image caption=Polonium in metallic form}}
|-
! [[Chemistry:Period 7 element|7]]
| {{element cell image|116|Livermorium|Lv| |Unknown phase|Other metal|Synthetic|legend=Other metal}}
|-
| colspan="2"|
----
''Legend''
{| style="text-align:center; border:0; margin: 0 auto"
|-
| style="border:{{element color|Primordial}}; background:{{Element color|table mark}};" | primordial element
|-
| style="border:{{element color|from decay}}; background:{{Element color|table mark}}; padding:0 2px;" | [[Physics:Radioactive decay|naturally occurring by radioactive decay]]
|-
| style="border:{{element color|Synthetic}}; background:{{Element color|table mark}};" | [[Physics:Synthetic element|synthetic element]]
|-
| Atomic number color:
|-
| <span style="color:{{element color|Gas}};">red=gas</span>,&nbsp;<span style="color:{{element color|Solid}};">black=solid</span>
|}
|}

The '''chalcogens''' (ore forming) ({{IPAc-en|ˈ|k|æ|l|k|ə|dʒ|ə|n|z}} {{Respell|KAL|kə|jənz}}) are the [[Chemistry:Chemical element|chemical element]]s in [[Chemistry:Group (periodic table)|group]] 16 of the [[Chemistry:Periodic table|periodic table]]. This group is also known as the '''oxygen family'''. Group 16 consists of the elements [[Chemistry:Oxygen|oxygen]] (O), [[Chemistry:Sulfur|sulfur]] (S), [[Chemistry:Selenium|selenium]] (Se), [[Chemistry:Tellurium|tellurium]] (Te), and the [[Physics:Radioactive decay|radioactive]] elements [[Chemistry:Polonium|polonium]] (Po) and [[Chemistry:Livermorium|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" is derived from a combination of the Greek word {{lang|grc-Latn|khalkόs}} ({{lang|grc|χαλκός}}) principally meaning [[Chemistry:Copper|copper]] (the term was also used for [[Chemistry:Bronze|bronze]]/[[Chemistry:Brass|brass]], any metal in the poetic sense, [[Earth:Ore|ore]] or [[Finance:Coin|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 the Latinized Greek word {{lang|la|genēs}}, meaning ''born'' or ''produced''.<ref>{{cite dictionary |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>

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 [[Physics:Valence electron|valence electron]]s, leaving them two electrons short of a full outer shell. Their most common [[Chemistry:Oxidation state|oxidation state]]s are −2, +2, +4, and +6. They have relatively low [[Physics:Atomic radius|atomic radii]], especially the lighter ones.<ref name="Jackson2002" />

All of the naturally occurring chalcogens have some role in biological functions, either as a nutrient or a toxin. Selenium is an important nutrient (among others as a building block of [[Chemistry:Selenocysteine|selenocysteine]]) but is also commonly toxic.<ref name="The Elements"/> Tellurium often has unpleasant effects (although some organisms can use it), and polonium (especially the [[Physics:Isotope|isotope]] [[Physics:Polonium-210|polonium-210]]) is always harmful as a result of its radioactivity.

Sulfur has more than 20 allotropes, oxygen has nine, selenium has at least eight, polonium has two, and only one crystal structure of tellurium has so far been discovered. There are numerous organic chalcogen compounds. Not counting oxygen, organic sulfur compounds are generally the most common, followed by organic selenium compounds and organic tellurium compounds. This trend also occurs with chalcogen [[Chemistry:Pnictogen|pnictides]] and compounds containing chalcogens and [[Chemistry:Carbon group|carbon group]] elements.

Oxygen is generally obtained by [[Physics:Air separation|separation]] of air into nitrogen and oxygen.<ref>{{Cite web |date=18 February 2021 |title=Saiba como é produzido o oxigênio hospitalar |url=https://cfq.org.br/noticia/saiba-como-e-produzido-o-oxigenio-hospitalar/ |access-date=December 23, 2023 |website=Conselho Federal de Química |archive-date=June 30, 2022 |archive-url=https://web.archive.org/web/20220630063513/http://cfq.org.br/noticia/saiba-como-e-produzido-o-oxigenio-hospitalar/ |url-status=dead }}</ref> Sulfur is extracted from oil and natural gas. Selenium and tellurium are produced as byproducts of copper refining. Polonium is most available in naturally occurring actinide-containing materials.  Livermorium has been synthesized in particle accelerators. The primary use of elemental oxygen is in [[Engineering:Steelmaking|steelmaking]].{{citation needed|date=July 2022}} Sulfur is mostly converted into [[Chemistry:Sulfuric acid|sulfuric acid]], which is heavily used in the chemical industry.<ref name="The Elements"/> Selenium's most common application is glassmaking. Tellurium compounds are mostly used in optical disks, electronic devices, and solar cells. Some of polonium's applications are due to its radioactivity.<ref name="ReferenceB"/>

==Properties==

===Atomic and physical===
Chalcogens show similar patterns in [[Physics:Electron configuration|electron configuration]], especially in the outermost [[Physics:Electron shell|shells]], where they all have the same number of [[Physics:Valence electron|valence electron]]s, resulting in similar trends in chemical behavior:

{| class="wikitable"
|-
![[Physics:Atomic number|Z]] !! [[Chemistry:Chemical element|Element]] !! [[Physics:Electron shell|No. of electrons/shell]]
|-
| 8 || Oxygen || 2, 6
|-
| 16 || Sulfur || 2, 8, 6
|-
| 34 || Selenium || 2, 8, 18, 6
|-
| 52 || Tellurium || 2, 8, 18, 18, 6
|-
| 84 || Polonium || 2, 8, 18, 32, 18, 6
|-
| 116 || Livermorium || 2, 8, 18, 32, 32, 18, 6 ''(predicted)''<ref name=Haire>{{Cite book |publisher= [[Company:Springer Science+Business Media|Springer Science+Business Media]] |location=Dordrecht, The Netherlands |bibcode=2011tcot.book.....M |doi =10.1007/978-94-007-0211-0 |editor1-last=Morss |editor1-first=Lester R. |editor2-last=Edelstein |editor2-first=Norman M. |editor3-last=Fuger| editor3-first=Jean |isbn =978-94-007-0210-3 |last1= Morss |first1= Lester R. |last2= Edelstein |first2= Norman M. |last3= Fuger |first3= Jean |year = 2011 |title = The Chemistry of the Actinide and Transactinide Elements }}</ref>
|}

{|class="wikitable" style="text-align:center"
|-
!Element !! Melting point<br />(°C)<ref name="Jackson2002" />
! Boiling point<br />(°C)<ref name="Jackson2002" />
!Density at STP<br />(g/cm<sup>3</sup>)<ref name="Jackson2002" />
|-
|Oxygen || −219 || −183 
|0.00143
|-
|Sulfur || 120 || 445 
|2.07
|-
|Selenium || 221 || 685 
|4.3
|-
|Tellurium || 450 || 988 
|6.24
|-
|Polonium || 254 || 962 
|9.2
|-
|Livermorium || 220 (predicted) || 800 (predicted) 
|14 (predicted)<ref name=Haire/>
|}

All chalcogens have six [[Physics:Valence electron|valence electron]]s. All of the solid, stable chalcogens are soft<ref>{{cite book|editor-last=Samsonov |editor-first=G. V. |chapter=Mechanical Properties of the Elements|doi=10.1007/978-1-4684-6066-7_7|isbn=978-1-4684-6066-7|chapter-url=http://ihtik.lib.ru/2011.08_ihtik_nauka-tehnika/2011.08_ihtik_nauka-tehnika_3560.rar|title=Handbook of the physicochemical properties of the elements|pages=387–446|publisher=IFI-Plenum|place=New York, USA|year=1968|url-status=dead|archive-url=https://web.archive.org/web/20150402123344/http://ihtik.lib.ru/2011.08_ihtik_nauka-tehnika/2011.08_ihtik_nauka-tehnika_3560.rar|archive-date=April 2, 2015|df=mdy-all}}</ref> and do not [[Physics:Thermal conductivity|conduct heat]] well.<ref name="Jackson2002" /> [[Chemistry:Electronegativity|Electronegativity]] decreases towards the chalcogens with higher atomic numbers. Density, melting and boiling points, and [[Physics:Atomic radius|atomic]] and [[Chemistry:Ionic radius|ionic radii]]<ref name="rsc">{{cite web |url=http://www.rsc.org/chemsoc/visualelements/pages/data/intro_groupvi_data.html |title=Visual Elements: Group 16 |publisher=Royal Society of Chemistry |access-date=November 25, 2013}}</ref> tend to increase towards the chalcogens with higher atomic numbers.<ref name="Jackson2002">{{cite book |last=Jackson |first=Mark |title=Periodic Table Advanced |publisher=Bar Charts Inc. |isbn=978-1-57222-542-8 |year=2002}}</ref>

===Isotopes===
Out of the six known chalcogens, one (oxygen) has an atomic number equal to a nuclear [[Physics:Magic number|magic number]], which means that their atomic nuclei tend to have increased stability towards radioactive decay.<ref name="The Disappearing Spoon"/> Oxygen has three theoretically stable isotopes, and 14 unstable isotopes. Sulfur has four theoretically stable isotopes, 20 radioactive isotopes, and one [[Physics:Nuclear isomer|isomer]]. Selenium has three theoretically stable isotopes, two observationally stable isotopes which are theoretically unstable, one nearly stable isotope, 26 radioactive isotopes, and 9 isomers. Tellurium has four theoretically stable isotopes, two observationally stable isotopes which are theoretically unstable, two nearly stable isotopes, 31 radioactive isotopes, and 17 isomers. Polonium has 42 isotopes, none of which are stable or nearly stable.<ref>{{cite web|last = Sonzogniurl|first = Alejandro|url = http://www.nndc.bnl.gov/nudat2/reCenter.jsp?z=84&n=130|title = Double Beta Decay for Selenium-82|access-date = November 25, 2013|publisher = Brookhaven National Laboratory|archive-date = October 3, 2021|archive-url = https://web.archive.org/web/20211003183335/https://www.nndc.bnl.gov/nudat2/reCenter.jsp?z=84&n=130|url-status = dead}}</ref> It has an additional 28 isomers.<ref name="ReferenceB"/> In addition to the stable isotopes, some radioactive chalcogen isotopes occur in nature, either because they are decay products, such as [[Physics:Polonium-210|<sup>210</sup>Po]], because they are [[Chemistry:Primordial nuclide|primordial]], such as <sup>82</sup>Se, because of [[Astronomy:Cosmic ray|cosmic ray]] [[Physics:Spallation|spallation]], or via [[Physics:Nuclear fission|nuclear fission]] of uranium. Livermorium isotopes <sup>290</sup>Lv through <sup>293</sup>Lv have been discovered; the most stable livermorium isotope is <sup>293</sup>Lv, which has a half-life of 0.061 seconds.<ref name = "ReferenceB"/><ref>{{cite journal|year = 1973|title = Double Beta Decay of Selenium-82|doi = 10.2113/gsecongeo.68.2.252|journal = Economic Geology|volume = 68|issue = 2|page = 252|last1 = Srinivasan|first1 = B.|last2 = Alexander|first2 = E. C.|last3 = Beaty|first3 = R. D.|last4 = Sinclair|first4 = D. E.|last5 = Manuel|first5 = O. K.}}</ref>

With the exception of oxygen and livermorium, all chalcogens have at least one naturally occurring radioisotope, sulfur has trace <sup>35</sup>S, selenium has <sup>82</sup>Se, tellurium has <sup>128</sup>Te and <sup>130</sup>Te, and polonium has <sup>210</sup>Po.

Among the lighter chalcogens (oxygen and sulfur), the most neutron-poor isotopes undergo [[Physics:Proton emission|proton emission]], the moderately neutron-poor isotopes undergo [[Physics:Electron capture|electron capture]] or β<sup>+</sup> decay, the moderately neutron-rich isotopes undergo [[Physics:Beta decay|β<sup>−</sup> decay]], and the most neutron rich isotopes undergo [[Physics:Neutron emission|neutron emission]]. The middle chalcogens (selenium and tellurium) have similar decay tendencies as the lighter chalcogens, but no proton-emitting isotopes have been observed, and some of the most neutron-deficient isotopes of tellurium undergo [[Physics:Alpha decay|alpha decay]]. Polonium isotopes tend to decay via alpha or beta decay.<ref>{{cite web |url=http://www.nndc.bnl.gov/nudat2/reCenter.jsp?z=47&n=63 |title=Nudat 2 |publisher=Nndc.bnl.gov |access-date=November 25, 2013 |archive-date=July 14, 2017 |archive-url=https://web.archive.org/web/20170714190234/http://www.nndc.bnl.gov/nudat2/reCenter.jsp?z=47&n=63 |url-status=dead }}</ref> Isotopes with nonzero [[Physics:Nuclear spin|nuclear spin]]s are more abundant in nature among the chalcogens selenium and tellurium than they are with sulfur.<ref name="synth"/>

===Allotropes===
[[File:Phase diagram of sulfur (1975).png|thumb|upright=1.2|Phase diagram of sulfur showing the relative stabilities of several allotropes<ref>{{cite web|url=https://www.osti.gov/biblio/4010212 |title=Phase Diagrams of the Elements|author=Young, David A. |date=September 11, 1975|publisher=Lawrence Livermore Laboratory |osti = 4010212}}</ref>]]
[[File:Chalkogene.jpg|thumb|left|The four stable chalcogens at [[Standard conditions for temperature and pressure|STP]]]]
[[File:Phase diagram of solid oxygen.svg|thumb|[[Phase diagram|Phase diagram]] for solid oxygen]]

Oxygen's most common allotrope is diatomic oxygen, or O<sub>2</sub>, a reactive paramagnetic molecule that is ubiquitous to [[Biology:Aerobic organism|aerobic organism]]s and has a blue color in its [[Physics:Liquid oxygen|liquid state]]. Another allotrope is O<sub>3</sub>, or [[Chemistry:Ozone|ozone]], which is three oxygen atoms bonded together in a bent formation. There is also an allotrope called [[Chemistry:Tetraoxygen|tetraoxygen]], or O<sub>4</sub>,<ref>{{cite journal|title = The ε Phase of Solid Oxygen: Evidence of an O4 Molecule Lattice|year = 1999|bibcode = 1999PhRvL..83.4093G|last1 = Gorelli|first1 = Federico A.|last2 = Ulivi|first2 = Lorenzo|last3 = Santoro|first3 = Mario|last4 = Bini|first4 = Roberto|volume = 83|page = 4093|journal = Physical Review Letters|doi = 10.1103/PhysRevLett.83.4093|issue = 20}}</ref> and six allotropes of [[Physics:Solid oxygen|solid oxygen]] including "red oxygen", which has the formula O<sub>8</sub>.<ref>{{cite journal|title = Observation of an O8 molecular lattice in the ε phase of solid oxygen|doi=10.1038/nature05174|journal =Nature|volume =443|issue =7108|pages =201–4|pmid =16971946|year = 2006|last1 = Lundegaard|first1 = Lars F.|last2 = Weck|first2 = Gunnar|last3 = McMahon|first3 = Malcolm I.|last4 = Desgreniers|first4 = Serge|last5 = Loubeyre|first5 = Paul|bibcode = 2006Natur.443..201L|s2cid=4384225}}</ref>

Sulfur has over 20 known allotropes, which is more than any other element except [[Chemistry: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 sulfur and [[Chemistry:Monoclinic|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&nbsp;°C, the chains begin to break down. If liquid sulfur above 190&nbsp;°C is [[Physics: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 [[Chemistry:Crystal structure|crystal structure]]. The gray allotrope of selenium is soft, with a [[Chemistry:Mohs hardness|Mohs hardness]] of 2, and brittle. Four other allotropes of selenium are metastable. These include two [[Chemistry:Monoclinic|monoclinic]] red allotropes and two [[Physics:Amorphous|amorphous]] allotropes, one of which is red and one of which is black.<ref>{{Cite web|vauthors = Butterman WC, ((Brown RD Jr)) |url =http://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|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&nbsp;°C.<ref name="ReferenceB"/>

The chalcogens have varying crystal structures. Oxygen's crystal structure is [[Physics:Monoclinic crystal system|monoclinic]], sulfur's is [[Physics:Orthorhombic crystal system|orthorhombic]], selenium and tellurium have the [[Physics:Hexagonal crystal system|hexagonal]] crystal structure, while polonium has a [[Cubic crystal system|cubic crystal structure]].<ref name="Jackson2002" /><ref name="The Elements">{{cite book|last = Gray|first = Theodore|title = The Elements|year = 2011|publisher = Black Bay and Leventhal publishers}}</ref>
{{clear}}

===Chemical===

Oxygen, sulfur, and selenium are [[Chemistry:Nonmetal|nonmetal]]s, and tellurium is a [[Chemistry:Metalloid|metalloid]], meaning that its chemical properties are between those of a [[Chemistry:Metal|metal]] and those of a nonmetal.<ref name="The Elements"/> It is not certain whether polonium is a metal or a metalloid. Some sources refer to polonium as a metalloid,<ref name="ReferenceB"/><ref name="Chemistry & reactivity">{{cite book|author1=Kotz, John C. |author2=Treichel, Paul M. |author3=Townsend, John Raymond |url = https://books.google.com/books?id=jcn6sgt7RpoC&pg=PA65 |title = Chemistry & Chemical Reactivity|year = 2009|page=65|publisher=Cengage Learning|isbn=978-0-495-38703-9}}</ref> although it has some metallic properties. Also, some allotropes of selenium display characteristics of a metalloid,<ref>{{cite web |url=http://www.gordonengland.co.uk/elements/metaloids.htm |title=Periodic Table of the Elements – Metalloids |publisher=Gordonengland.co.uk |access-date=November 25, 2013}}</ref> even though selenium is usually considered a nonmetal. Even though oxygen is a chalcogen, its chemical properties are different from those of other chalcogens. One reason for this is that the heavier chalcogens have vacant d-orbitals. Oxygen's electronegativity is also much higher than those of the other chalcogens. This makes oxygen's electric polarizability several times lower than those of the other chalcogens.<ref name="synth">{{cite book |author = Zakai, Uzma I. |url = https://books.google.com/books?id=k-LjiXfTXnYC |title = Design, Synthesis, and Evaluation of Chalcogen Interactions |year = 2007 |isbn = 978-0-549-34696-8 |access-date = November 25, 2013 }}{{Dead link|date=November 2023 |bot=InternetArchiveBot |fix-attempted=yes }}</ref>

For [[Chemistry:Covalent bond|covalent bond]]ing a chalcogen may accept two electrons according to the [[Octet rule|octet rule]], leaving two [[Chemistry:Lone pair|lone pair]]s. When an atom forms two [[Chemistry:Single bond|single bond]]s, they [[Physics:Bent molecular geometry|form an angle between 90° and 120°]]. In 1+ cations, such as {{chem2|H3O+}}, a chalcogen forms three [[Chemistry:Molecular orbital|molecular orbital]]s arranged in a [[Physics:Trigonal pyramidal molecular geometry|trigonal pyramidal]] fashion and one lone pair. Double bonds are also common in chalcogen compounds, for example in chalcogenates (see below).

The oxidation number of the most common chalcogen compounds with positive metals is −2. However the tendency for chalcogens to form compounds in the −2 state decreases towards the heavier chalcogens.<ref name="wisc"/> Other oxidation numbers, such as −1 in [[Chemistry:Pyrite|pyrite]] and [[Chemistry:Peroxide|peroxide]], do occur. The highest formal oxidation number is +6.<ref name="Jackson2002" /> This oxidation number is found in [[Chemistry:Sulfate|sulfate]]s, [[Chemistry:Selenate|selenate]]s, [[Chemistry:Tellurate|tellurate]]s, polonates, and their corresponding acids, such as [[Chemistry:Sulfuric acid|sulfuric acid]].

Oxygen is the most electronegative element except for [[Chemistry:Fluorine|fluorine]], and forms compounds with almost all of the chemical elements, including some of the [[Chemistry:Noble gas|noble gas]]es. It commonly bonds with many metals and metalloids to form [[Chemistry:Oxide|oxide]]s, including [[Chemistry:Iron oxide|iron oxide]], [[Chemistry:Titanium oxide|titanium oxide]], and silicon oxide. Oxygen's most common [[Chemistry:Oxidation state|oxidation state]] is −2, and the oxidation state −1 is also relatively common.<ref name="Jackson2002" /> With [[Chemistry:Hydrogen|hydrogen]] it forms water and [[Chemistry:Hydrogen peroxide|hydrogen peroxide]]. Organic oxygen compounds are ubiquitous in [[Chemistry:Organic chemistry|organic chemistry]]. <!--the text is broken apart too much, they're having five stories and not one-->

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 [[Chemistry:Double bond|double bond]]s are far weaker than oxygen-oxygen double bonds, but sulfur-sulfur [[Chemistry:Single bond|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 [[Chemistry:Thiol|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 [[Chemistry:Organoselenium chemistry|organic selenium compounds]], such as selenoproteins. Tellurium's oxidation states are −2, +2, +4, and +6.<ref name="Jackson2002" /> Tellurium forms the oxides [[Chemistry:Tellurium monoxide|tellurium monoxide]], [[Chemistry:Tellurium dioxide|tellurium dioxide]], and [[Chemistry:Tellurium trioxide|tellurium trioxide]].<ref name="ReferenceB"/> Polonium's oxidation states are +2 and +4.<ref name="Jackson2002" />
[[File:Brindis (24675281395).jpg|thumb|upright|left|[[Chemistry:Water|Water]] ({{chem2|H2O}}) is the most familiar chalcogen-containing compound.|alt=Water dripping into a glass, showing drops and bubbles.]]

There are many acids containing chalcogens, including sulfuric acid, [[Chemistry:Sulfurous acid|sulfurous acid]], [[Chemistry:Selenic acid|selenic acid]], and [[Chemistry:Telluric acid|telluric acid]]. All [[Chemistry:Hydrogen chalcogenide|hydrogen chalcogenide]]s are toxic except for [[Chemistry:Water|water]].<ref name="Emsley2011">{{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><ref>{{cite web |last1 = Van Vleet|first1 = JF|last2=Boon |first2=GD |last3=Ferrans |first3=VJ|url = http://toxnet.nlm.nih.gov/cgi-bin/sis/search/a?dbs+hsdb:@term+@DOCNO+7057|title = Tellurium compounds|year = 1981|publisher=The Toxicology and Environmental Health Information Program, US National Institutes of Health |access-date=November 25, 2013}}</ref> Oxygen ions often come in the forms of [[Chemistry:Oxide|oxide]] ions ({{chem2|O(2-)}}), [[Chemistry:Peroxide|peroxide]] ions ({{chem2|O2(2-)}}), and [[Chemistry:Hydroxide|hydroxide]] ions ({{chem2|OH-}}). Sulfur ions generally come in the form of [[Chemistry:Sulfide|sulfide]]s ({{chem2|S(2-)}}), [[Chemistry:Bisulfide|bisulfide]]s ({{chem2|SH-}}), [[Chemistry:Sulfite|sulfite]]s ({{chem2|SO3(2-)}}), [[Chemistry:Sulfate|sulfate]]s ({{chem2|SO4(2-)}}), and [[Chemistry:Thiosulfate|thiosulfate]]s ({{chem2|S2O3(2-)}}). Selenium ions usually come in the form of [[Physics:Selenide|selenide]]s ({{chem2|Se(2-)}}), [[Chemistry:Selenite (ion)|selenite]]s ({{chem2|SeO3(2-)}}) and [[Chemistry:Selenate|selenate]]s ({{chem2|SeO4(2-)}}). Tellurium ions often come in the form of [[Chemistry:Tellurate|tellurate]]s ({{chem2|TeO4(2-)}}).<ref name="Jackson2002" /> Molecules containing metal bonded to chalcogens are common as minerals. For example, [[Chemistry:Pyrite|pyrite]] (FeS<sub>2</sub>) is an [[Earth:Iron ore|iron ore]], and the rare mineral [[Chemistry:Calaverite|calaverite]] is the ditelluride {{chem2|(Au, Ag)Te2}}.

Although all group 16 elements of the periodic table, including oxygen, can be defined as chalcogens, oxygen and oxides are usually distinguished from chalcogens and chalcogenides. The term ''chalcogenide'' is more commonly reserved for [[Chemistry:Sulfide|sulfide]]s, [[Physics:Selenide|selenide]]s, and [[Chemistry:Telluride|telluride]]s, rather than for [[Chemistry:Oxide|oxide]]s<!--then why are you talking about oxygen here at all?-->.<ref name="chalcogen2" /><ref name='handbook'>{{cite book |url=https://books.google.com/books?id=IvGnUAaSqOsC&pg=PT24 |editor1=Devillanova, Francesco |title=Handbook of Chalcogen Chemistry –New Perspectives in Sulfur, Selenium and Tellurium |publisher=Royal Society of Chemistry |year=2007 |isbn=978-0-85404-366-8 |access-date=November 25, 2013}}</ref><ref name='Takahisa'>{{cite journal|doi=10.1016/0039-6028(91)90679-M|title=Passivation of GaAs(001) surfaces by chalcogen atoms (S, Se and Te)|year=1991|last1=Takahisa|first1=Ohno|journal=Surface Science|volume=255|issue=3|page=229|bibcode = 1991SurSc.255..229T }}</ref>

Except for polonium, the chalcogens are all fairly similar to each other chemically. They all form X<sup>2−</sup> ions when reacting with electropositive metals.<ref name="wisc">{{cite web|url=http://chemed.chem.wisc.edu/chempaths/GenChem-Textbook/Group-VIA-Chalcogens-609.html |title=Group VIA: Chalcogens |publisher=Chemed.chem.wisc.edu |url-status=dead |access-date=November 25, 2013 |archive-url=https://web.archive.org/web/20131104164205/http://chemed.chem.wisc.edu/chempaths/GenChem-Textbook/Group-VIA-Chalcogens-609.html |archive-date=November 4, 2013 }}</ref>

[[Chemistry:Sulfide mineral|Sulfide mineral]]s and analogous compounds produce gases upon reaction with oxygen.<ref>{{cite journal |url=http://www.minersoc.org/pages/Archive-MM/Volume_57/57-389-599.pdf |archive-url=https://web.archive.org/web/20131029185906/http://www.minersoc.org/pages/Archive-MM/Volume_57/57-389-599.pdf |archive-date=2013-10-29 |url-status=live |title=Mineral deposits and chalcogen gases |author=Hale, Martin |journal=Mineralogical Magazine |year=1993 |volume=57 |pages=599–606|doi=10.1180/minmag.1993.057.389.04|issue=389 |access-date=November 25, 2013|bibcode=1993MinM...57..599H|citeseerx=10.1.1.606.8357 }}</ref>
{{clear}}

==Compounds==

===With halogens=== <!-- Chalcohalide redirects here -->
Chalcogens also form compounds with [[Chemistry:Halogen|halogen]]s known as '''chalcohalides''', or '''chalcogen halides'''. The majority of simple chalcogen halides are well-known and widely used as chemical [[Biology:Reagent|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 [[Chemistry:Sulfur hexafluoride|sulfur hexafluoride]], are stable and well-known. [[Chemistry:Sulfur tetrafluoride|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 [[Chemistry:Selenium tetrafluoride|selenium tetrafluoride]] and [[Chemistry:Tellurium tetrafluoride|tellurium tetrafluoride]] are known. Chalcogen chlorides and bromides have also been explored. In particular, selenium dichloride and sulfur dichloride can react to form [[Chemistry: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.{{dubious|date=September 2014}} Such compounds can form in mixtures of [[Chemistry:Sulfur dichloride|sulfur dichloride]] and selenium halides. 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|MoS2Cl3}}. Unlike selenium chlorides and bromides, selenium [[Chemistry:Iodide|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 [[Chemistry:Polymer|polymer]]s when in the [[Chemistry: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 anions.<ref name="handbook"/>

===Organic===
[[Chemistry:Alcohol|Alcohol]]s, [[Chemistry:Phenol|phenol]]s and other similar compounds contain oxygen. However, in [[Chemistry:Thiol|thiol]]s, [[Chemistry:Selenol|selenol]]s and [[Chemistry:Tellurol|tellurol]]s; sulfur, selenium, and tellurium replace oxygen. Thiols are better known than selenols or tellurols. Aside from alcohols, thiols are the most stable chalcogenols and tellurols are the least stable, being unstable in heat or light. Other organic chalcogen compounds include [[Chemistry:Thioether|thioether]]s, selenoethers and telluroethers. Some of these, such as [[Chemistry:Dimethyl sulfide|dimethyl sulfide]], [[Chemistry:Diethyl sulfide|diethyl sulfide]], and dipropyl sulfide are commercially available. Selenoethers are in the form of [[Chemistry:Side chain|R]]<sub>2</sub>Se or [[Chemistry:Side chain|R]]SeR. Telluroethers such as [[Chemistry:Dimethyl telluride|dimethyl telluride]] are typically prepared in the same way as thioethers and selenoethers. Organic chalcogen compounds, especially organic sulfur compounds, have the tendency to smell unpleasant. Dimethyl telluride also smells unpleasant,<ref>{{cite encyclopedia |url=https://www.britannica.com/EBchecked/topic/592252/thiol |title=thiol (chemical compound)|encyclopedia = [[Encyclopædia Britannica]] |access-date=November 25, 2013}}</ref> and selenophenol is renowned for its "metaphysical stench".<ref>{{cite web|vauthors = Lowe D|url=http://pipeline.corante.com/archives/2012/05/15/things_i_wont_work_with_selenophenol.php |title=Things I Won't Work With: Selenophenol |work=In the Pipeline |date=May 15, 2012 |access-date=November 25, 2013 |archive-url=https://web.archive.org/web/20120515170845/http://pipeline.corante.com/archives/2012/05/15/things_i_wont_work_with_selenophenol.php |archive-date=May 15, 2012 |df=mdy }}</ref> There are also [[Chemistry:Thioketone|thioketone]]s, selenoketones, and telluroketones. Out of these, thioketones are the most well-studied with 80% of chalcogenoketones papers being about them. Selenoketones make up 16% of such papers and telluroketones make up 4% of them. Thioketones have well-studied non-linear electric and photophysical properties. Selenoketones are less stable than thioketones and telluroketones are less stable than selenoketones. Telluroketones have the highest level of [[Chemistry:Chemical polarity|polarity]] of chalcogenoketones.<ref name="handbook"/>

===With metals===

There is a very large number of metal chalcogenides. There are also ternary compounds containing [[Chemistry:Alkali metal|alkali metal]]s and [[Chemistry:Transition metal|transition metal]]s. Highly metal-rich metal chalcogenides, such as [[Chemistry:Lutetium|Lu]]<sub>7</sub>Te and Lu<sub>8</sub>Te have domains of the metal's crystal lattice containing chalcogen atoms. While these compounds do exist, analogous chemicals that contain [[Chemistry:Lanthanum|lanthanum]], [[Chemistry:Praseodymium|praseodymium]], [[Chemistry:Gadolinium|gadolinium]], [[Chemistry:Holmium|holmium]], [[Chemistry:Terbium|terbium]], or [[Chemistry:Ytterbium|ytterbium]] have not been discovered, as of 2008. The [[Chemistry:Boron group|boron group]] metals aluminum, [[Chemistry:Gallium|gallium]], and [[Chemistry:Indium|indium]] also form bonds to chalcogens. The Ti<sup>3+</sup> ion forms chalcogenide [[Chemistry:Dimer|dimers]] such as Ti[[Chemistry:Thallium|Tl]]<sub>5</sub>Se<sub>8</sub>. Metal chalcogenide dimers also occur as lower tellurides, such as Zr<sub>5</sub>Te<sub>6</sub>.<ref name="handbook"/>

Elemental chalcogens react with certain lanthanide compounds to form lanthanide clusters rich in chalcogens.{{dubious|date=September 2014}} [[Chemistry:Uranium|Uranium]](IV) chalcogenol compounds also exist. There are also [[Chemistry:Transition metal|transition metal]] chalcogenols which have potential to serve as [[Chemistry:Catalyst|catalyst]]s and stabilize [[Physics:Nanoparticle|nanoparticle]]s.<ref name="handbook"/>

===With pnictogens===

[[File:Sulfid bismutitý.PNG|thumb|Bismuth sulfide, a pnictogen chalcogenide]]
Compounds with chalcogen-[[Chemistry:Phosphorus|phosphorus]] bonds have been explored for more than 200 years. These compounds include unsophisticated phosphorus chalcogenides as well as large molecules with biological roles and phosphorus-chalcogen compounds with metal clusters. These compounds have numerous applications, including organo-phosphate insecticides, [[Engineering:Match#The strike-anywhere match|strike-anywhere matches]] and [[Physics:Quantum dot|quantum dot]]s. A total of 130,000 compounds with at least one phosphorus-sulfur bond, 6000 compounds with at least one phosphorus-selenium bond, and 350 compounds with at least one phosphorus-tellurium bond have been discovered.{{citation needed|date=September 2014}} The decrease in the number of chalcogen-phosphorus compounds further down the periodic table is due to diminishing bond strength. Such compounds tend to have at least one phosphorus atom in the center, surrounded by four chalcogens and [[Chemistry:Side chain|side chain]]s. However, some phosphorus-chalcogen compounds also contain hydrogen (such as secondary [[Chemistry:Phosphine|phosphine]] chalcogenides) or nitrogen (such as dichalcogenoimidodiphosphates). [[Chemistry:Phosphorus selenide|Phosphorus selenide]]s are typically harder to handle that phosphorus sulfides, and compounds in the form P<sub>x</sub>Te<sub>y</sub> have not been discovered. Chalcogens also bond with other pnictogens, such as [[Chemistry:Arsenic|arsenic]], [[Chemistry:Antimony|antimony]], and [[Chemistry:Bismuth|bismuth]]. Heavier chalcogen pnictides tend to form [[Engineering:Ribbon|ribbon]]-like polymers instead of individual molecules. Chemical formulas of these compounds include Bi<sub>2</sub>S<sub>3</sub> and Sb<sub>2</sub>Se<sub>3</sub>. Ternary chalcogen pnictides are also known. Examples of these include P<sub>4</sub>O<sub>6</sub>Se and P<sub>3</sub>SbS<sub>3</sub>. [[Chemistry:Salt|salts]] containing chalcogens and pnictogens also exist. Almost all chalcogen pnictide salts are typically in the form of [Pn<sub>x</sub>E<sub>4x</sub>]<sup>3−</sup>, where Pn is a pnictogen and E is a chalcogen.{{dubious|date=September 2014}} Tertiary phosphines can react with chalcogens to form compounds in the form of R<sub>3</sub>PE, where E is a chalcogen. When E is sulfur, these compounds are relatively stable, but they are less so when E is selenium or tellurium. Similarly, secondary phosphines can react with chalcogens to form secondary phosphine chalcogenides. However, these compounds are in a state of [[Chemistry:Chemical equilibrium|equilibrium]] with chalcogenophosphinous acid. Secondary phosphine chalcogenides are weak acids.<ref name="handbook"/> Binary compounds consisting of antimony or arsenic and a chalcogen. These compounds tend to be colorful and can be created by a reaction of the constituent elements at temperatures of {{convert|500|to|900|C|F}}.<ref>{{Citation|author1=A. Earnshaw |author2=Norman Greenwood |url = https://books.google.com/books?id=EvTI-ouH3SsC&q=chalcogen|title = Chemistry of the Elements|date = November 11, 1997|publisher=Elsevier |access-date = February 12, 2014|isbn = 9780080501093}}</ref>

===Other===
Chalcogens form single bonds and double bonds with other [[Chemistry:Carbon group|carbon group]] elements than carbon, such as [[Chemistry:Silicon|silicon]], [[Chemistry:Germanium|germanium]], and [[Chemistry:Tin|tin]]. Such compounds typically form from a reaction of carbon group halides and chalcogenol salts or chalcogenol [[Chemistry:Base|bases]]. Cyclic compounds with chalcogens, carbon group elements, and boron atoms exist, and occur from the reaction of boron dichalcogenates and carbon group metal halides. Compounds in the form of M-E, where M is silicon, germanium, or tin, and E is sulfur, selenium or tellurium have been discovered. These form when carbon group [[Physics:Hydride|hydride]]s react or when heavier versions of [[Chemistry:Carbene|carbene]]s react.{{dubious|date=September 2014}} Sulfur and tellurium can bond with organic compounds containing both silicon and phosphorus.<ref name="handbook"/>

All of the chalcogens form [[Physics:Hydride|hydride]]s. In some cases this occurs with chalcogens bonding with two hydrogen atoms.<ref name="ReferenceB"/> However tellurium hydride and [[Chemistry:Polonium hydride|polonium hydride]] are both volatile and highly labile.<ref name = "holleman">{{cite book|editor1=Holleman, Arnold F. |editor2=Wiber, Egon |editor3=Wiberg, Nils |url = https://books.google.com/books?id=Mtth5g59dEIC&pg=PA470|title = Inorganic Chemistry|pages=470 ff|isbn = 978-0-12-352651-9|year=2001|publisher=Academic Press }}</ref> Also, oxygen can bond to hydrogen in a 1:1 ratio as in [[Chemistry:Hydrogen peroxide|hydrogen peroxide]], but this compound is unstable.<ref name="wisc"/>

Chalcogen compounds form a number of [[Chemistry:Interchalcogen|interchalcogen]]s. For instance, sulfur forms the toxic [[Chemistry:Sulfur dioxide|sulfur dioxide]] and [[Chemistry:Sulfur trioxide|sulfur trioxide]].<ref name="wisc"/> Tellurium also forms oxides. There are some chalcogen sulfides as well. These include selenium sulfide, an ingredient in some [[Medicine:Shampoo|shampoo]]s.<ref name="The Elements"/>

Since 1990, a number of [[Chemistry:Boride|boride]]s with chalcogens bonded to them have been detected. The chalcogens in these compounds are mostly sulfur, although some do contain selenium instead. One such chalcogen boride consists of two molecules of [[Chemistry:Dimethyl sulfide|dimethyl sulfide]] attached to a boron-hydrogen molecule. Other important boron-chalcogen compounds include macropolyhedral systems. Such compounds tend to feature sulfur as the chalcogen. There are also chalcogen borides with two, three, or four chalcogens. Many of these contain sulfur but some, such as Na<sub>2</sub>B<sub>2</sub>Se<sub>7</sub> contain selenium instead.<ref>{{cite book|editor = Devillanova, Francesco A. |title=Handbook of chalcogen chemistry|url=https://books.google.com/books?id=IvGnUAaSqOsC|year = 2007|publisher=Royal Society of Chemistry |isbn = 978-0-85404-366-8 |access-date=November 25, 2013}}</ref>

==History==

===Early discoveries===
<!--Not supported by its article: [[File:Greekfire-madridskylitzes1.jpg|thumb|Greek fire, an early sulfur-related discovery]]-->
Sulfur has been known since ancient times and is mentioned in the [[Religion:Bible|Bible]] fifteen times. It was known to the ancient Greeks and commonly mined by the ancient Romans. <!--Not supported by the article: It was also historically used as a component of [[Engineering:Greek fire|Greek fire]].--> In the Middle Ages, it was a key part of [[Philosophy:Alchemy|alchemical]] experiments. In the 1700s and 1800s, scientists [[Biography:Joseph Louis Gay-Lussac|Joseph Louis Gay-Lussac]] and Louis-Jacques Thénard proved sulfur to be a chemical element.<ref name="ReferenceB"/>

Early attempts to separate oxygen from air were hampered by the fact that air was thought of as a single element up to the 17th and 18th centuries. Robert Hooke, [[Biography:Mikhail Lomonosov|Mikhail Lomonosov]], Ole Borch, and Pierre Bayden all successfully created oxygen, but did not realize it at the time. Oxygen was discovered by [[Biography:Joseph Priestley|Joseph Priestley]] in 1774 when he focused sunlight on a sample of mercuric oxide and collected the resulting gas. [[Biography:Carl Wilhelm Scheele|Carl Wilhelm Scheele]] had also created oxygen in 1771 by the same method, but Scheele did not publish his results until 1777.<ref name="ReferenceB"/>

Tellurium was first discovered in 1783 by Franz Joseph Müller von Reichenstein. He discovered tellurium in a sample of what is now known as calaverite. Müller assumed at first that the sample was pure antimony, but tests he ran on the sample did not agree with this. Muller then guessed that the sample was bismuth sulfide, but tests confirmed that the sample was not that. For some years, Muller pondered the problem. Eventually he realized that the sample was gold bonded with an unknown element. In 1796, Müller sent part of the sample to the German chemist Martin Klaproth, who purified the undiscovered element. Klaproth decided to call the element tellurium after the Latin word for earth.<ref name="ReferenceB"/>

Selenium was discovered in 1817 by [[Biography:Jöns Jacob Berzelius|Jöns Jacob Berzelius]]. Berzelius noticed a reddish-brown sediment at a sulfuric acid manufacturing plant. The sample was thought to contain arsenic. Berzelius initially thought that the sediment contained tellurium, but came to realize that it also contained a new element, which he named selenium after the Greek moon goddess Selene.<ref name="ReferenceB"/><ref>{{cite journal|last = Trofast|first = Jan|url = http://www.iupac.org/publications/ci/2011/3305/5_trofast.html|title = Berzelius' Discovery of Selenium|journal= Chemistry International|volume=33 |issue=5|date=September–October 2011 |access-date=November 25, 2013}}</ref>

===Periodic table placing===
[[File:Mendelejevs periodiska system 1871.png|thumb|upright=1.5|Dmitri Mendeleev's periodic system proposed in 1871 showing oxygen, sulfur, selenium and tellurium part of his group VI]]

Three of the chalcogens (sulfur, selenium, and tellurium) were part of the discovery of [[Chemistry:Periodic table|periodicity]], as they are among a series of triads of elements in the same [[Chemistry:Group (periodic table)|group]] that were noted by Johann Wolfgang Döbereiner as having similar properties.<ref name="The Disappearing Spoon"/> Around 1865 John Newlands produced a series of papers where he listed the elements in order of increasing atomic weight and similar physical and chemical properties that recurred at intervals of eight; he likened such periodicity to the octaves of music.<ref>{{Cite journal |title = On Relations Among the Equivalents |author = Newlands, John A. R. |journal = Chemical News |volume = 10 |pages = 94–95 |date =20 August 1864 |url=http://web.lemoyne.edu/~GIUNTA/EA/NEWLANDSann.HTML |url-status=live |archive-url=https://web.archive.org/web/20110101073248/http://web.lemoyne.edu/~GIUNTA/EA/NEWLANDSann.HTML |archive-date=January 1, 2011 |access-date=November 25, 2013}}</ref><ref>{{cite journal |title = On the Law of Octaves |author = Newlands, John A. R. |journal = Chemical News |volume = 12 |page = 83 |date = 18 August 1865 |url=http://web.lemoyne.edu/~GIUNTA/EA/NEWLANDSann.HTML |url-status=live |archive-url=https://web.archive.org/web/20110101073248/http://web.lemoyne.edu/~GIUNTA/EA/NEWLANDSann.HTML |archive-date=January 1, 2011 |access-date=November 25, 2013}}</ref> His version included a "group b" consisting of oxygen, sulfur, selenium, tellurium, and [[Chemistry:Osmium|osmium]].

[[File:Johann Wolfgang Döbereiner.jpg|thumb|upright|left|Johann Wolfgang Döbereiner was among the first to notice similarities between what are now known as chalcogens.]]

After 1869, [[Biography:Dmitri Mendeleev|Dmitri Mendeleev]] proposed his periodic table placing oxygen at the top of "group VI" above sulfur, selenium, and tellurium.<ref>{{cite journal|last=Mendelejew |first=Dimitri |year=1869 |title=Über die Beziehungen der Eigenschaften zu den Atomgewichten der Elemente |journal=Zeitschrift für Chemie |pages=405–406 |language=de}}</ref> [[Chemistry:Chromium|Chromium]], [[Chemistry:Molybdenum|molybdenum]], [[Chemistry:Tungsten|tungsten]], and [[Chemistry:Uranium|uranium]] were sometimes included in this group, but they would be later rearranged as part of [[Chemistry:Group 6 element|group VIB]]; uranium would later be moved to the [[Chemistry:Actinide|actinide]] series. Oxygen, along with sulfur, selenium, tellurium, and later polonium would be grouped in ''group VIA'', until the group's name was changed to ''group 16'' in 1988.<ref>{{cite journal |last1=Fluck |first1=E.|year=1988 |title=New Notations in the Periodic Table |journal=[[Chemistry:Pure and Applied Chemistry|Pure Appl. Chem.]]|volume=60|pages=431–436|doi=10.1351/pac198860030431|url=http://www.iupac.org/publications/pac/1988/pdf/6003x0431.pdf |archive-url=https://ghostarchive.org/archive/20221010/http://www.iupac.org/publications/pac/1988/pdf/6003x0431.pdf |archive-date=2022-10-10 |url-status=live|issue=3 |s2cid=96704008|access-date=November 25, 2013}}</ref>

===Modern discoveries===
In the late 19th century, [[Biography:Marie Curie|Marie Curie]] and [[Biography:Pierre Curie|Pierre Curie]] discovered that a sample of pitchblende was emitting four times as much radioactivity as could be explained by the presence of uranium alone. The Curies gathered several tons of pitchblende and refined it for several months until they had a pure sample of polonium. The discovery officially took place in 1898. Prior to the invention of particle accelerators, the only way to produce polonium was to extract it over several months from uranium ore.<ref name="ReferenceB"/>

The first attempt at creating livermorium was from 1976 to 1977 at the LBNL, who bombarded curium-248 with calcium-48, but were not successful. After several failed attempts in 1977, 1998, and 1999 by research groups in Russia, Germany, and the US, livermorium was created successfully in 2000 at the [[Organization:Joint Institute for Nuclear Research|Joint Institute for Nuclear Research]] by bombarding [[Chemistry:Curium|curium]]-248 atoms with calcium-48 atoms. The element was known as ununhexium until it was officially named livermorium in 2012.<ref name="ReferenceB"/>

===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|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 [[Chemistry:Oxyacid|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}}, literally "[[Chemistry:Copper|copper]]"), 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). It was first used in 1932 by Wilhelm Biltz's group at Leibniz University Hannover, where it was proposed by Werner Fischer.<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 [[Chemistry:Bronze|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 [[Social:Sanskrit|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 [[Organization:Lawrence Livermore National Laboratory|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>

==Occurrence==
The four lightest chalcogens (oxygen, sulfur, selenium, and tellurium) are all primordial elements on Earth. Sulfur and oxygen occur as constituent copper ores and selenium and tellurium occur in small traces in such ores.<ref name="wisc"/> [[Chemistry:Polonium|Polonium]] forms naturally from the decay of other elements, even though it is not primordial. Livermorium does not occur naturally at all.

Oxygen makes up 21% of the atmosphere by weight, 89% of water by weight, 46% of the Earth's crust by weight,<ref name="Jackson2002" /> and 65% of the human body.<ref name="Structure of Matter">{{cite book|last = Galan|first = Mark|title = Structure of Matter|year = 1992|publisher = International Editorial Services Inc.|isbn=978-0-8094-9662-4}}</ref> Oxygen also occurs in many minerals, being found in all [[Chemistry:Oxide minerals|oxide minerals]] and hydroxide minerals, and in numerous other mineral groups.<ref name="Rocks and Minerals">{{cite book|last = Pellant|first = Chris|title = Rocks and Minerals|year = 1992|publisher = Dorling Kindserley handbooks|isbn=978-0-7513-2741-0}}</ref> Stars of at least eight times the mass of the Sun also produce oxygen in their cores via [[Physics:Nuclear fusion|nuclear fusion]].<ref name="The Disappearing Spoon">{{cite book|last = Kean|first = Sam|title = The Disappearing Spoon|year = 2011|publisher = Back Bay Books|isbn=978-0-316-05163-7}}</ref> Oxygen is the third-most [[Chemistry:Abundance of the chemical elements#Universe|abundant element in the universe]], making up 1% of the universe by weight.<ref name = "Glossary">{{cite web|last=Heiserman|first=Davis L.|title=The 10 Most Abundant Elements in the Universe|url=http://education.jlab.org/glossary/abund_uni.html|year = 1992|access-date = February 6, 2013}}</ref><ref name = "Abundance"/>

Sulfur makes up 0.035% of the Earth's crust by weight, making it the 17th most abundant element there<ref name="Jackson2002" /> and makes up 0.25% of the human body.<ref name="Structure of Matter"/> It is a major component of soil. Sulfur makes up 870 parts per million of seawater and about 1 part per billion of the atmosphere.<ref name="ReferenceB"/> Sulfur can be found in elemental form or in the form of [[Chemistry:Sulfide minerals|sulfide minerals]], [[Chemistry:Sulfate minerals|sulfate minerals]], or [[Chemistry:Sulfosalt minerals|sulfosalt minerals]].<ref name="Rocks and Minerals"/> Stars of at least 12 times the mass of the Sun produce sulfur in their cores via nuclear fusion.<ref name="The Disappearing Spoon"/> Sulfur is the tenth most abundant element in the universe, making up 500 parts per million of the universe by weight.<ref name="Glossary"/><ref name="Abundance">{{cite web|last = Winter|first = Mark|title = Abundance in the universe|url = http://www.webelements.com/periodicity/abundance_universe/|year = 1993|access-date = February 6, 2013|archive-url = https://web.archive.org/web/20130117100231/http://www.webelements.com/periodicity/abundance_universe/|archive-date = January 17, 2013|url-status = dead|df = mdy-all}}</ref>

Selenium makes up 0.05 [[Parts per million|parts per million]] of the Earth's crust by weight.<ref name="Jackson2002" /> This makes it the 67th most [[Chemistry:Abundance of the chemical elements#Earth|abundant element in the Earth's crust]]. Selenium makes up on average 5 parts per million of the soils. [[Earth:Seawater|Seawater]] contains around 200 parts per trillion of selenium. The atmosphere contains 1 nanogram of selenium per cubic meter. There are mineral groups known as selenates and selenites, but there are not many minerals in these groups.<ref>{{cite web|author = Amethyst Galleries |url = http://www.galleries.com/minerals/sulfates/subclass.htm#selenates|title = Sulfates|year = 1995 |access-date=November 25, 2013}}</ref> Selenium is not produced directly by nuclear fusion.<ref name = "The Disappearing Spoon"/> Selenium makes up 30 parts per billion of the universe by weight.<ref name="Abundance"/>

There are only 5 parts per billion of tellurium in the Earth's crust and 15 parts per billion of tellurium in seawater.<ref name="ReferenceB"/> Tellurium is one of the eight or nine least abundant elements in the Earth's crust.<ref name="The Elements"/> There are a few dozen tellurate minerals and telluride minerals, and tellurium occurs in some minerals with gold, such as [[Chemistry:Sylvanite|sylvanite]] and calaverite.<ref>{{cite web|author = Amethyst Galleries |url = http://www.galleries.com/minerals/sulfates/subclass.htm#tellurates|title = Tellurates|year = 1995 |access-date=November 25, 2013}}</ref> Tellurium makes up 9 parts per billion of the universe by weight.<ref name="The Elements"/><ref name="Abundance"/><ref>{{cite web|author = Advameg |website=Chemistry Explained |url = http://www.chemistryexplained.com/elements/P-T/Tellurium.html|title = Tellurium|year = 2013 |access-date=November 25, 2013}}</ref>

Polonium only occurs in trace amounts on Earth, via radioactive decay of uranium and thorium. It is present in uranium ores in concentrations of 100 micrograms per metric ton. Very minute amounts of polonium exist in the soil and thus in most food, and thus in the human body.<ref name="ReferenceB"/> The Earth's crust contains less than 1 part per billion of polonium, making it one of the ten rarest metals on Earth.<ref name="ReferenceB"/><ref name="Jackson2002" />

Livermorium is always produced artificially in [[Physics:Particle accelerator|particle accelerator]]s. Even when it is produced, only a small number of atoms are synthesized at a time.

===Chalcophile elements===
Chalcophile elements are those that remain on or close to the surface because they combine readily with chalcogens other than oxygen, forming compounds which do not sink into the core. Chalcophile ("chalcogen-loving") elements in this context are those metals and heavier nonmetals that have a low affinity for oxygen and prefer to bond with the heavier chalcogen sulfur as sulfides.<ref>{{cite book|last = Sodhi|first = G. S.|url = https://books.google.com/books?id=EMlxjftDCY0C&pg=PA33|title = Fundamental Concepts of Environmental Chemistry|year = 2000| publisher=Alpha Science International |isbn = 978-1-84265-281-7}}</ref> Because sulfide minerals are much denser than the silicate minerals formed by lithophile elements,<ref name = "Rocks and Minerals"/> chalcophile elements separated below the lithophiles at the time of the first crystallisation of the Earth's crust. This has led to their depletion in the Earth's crust relative to their solar abundances, though this depletion has not reached the levels found with siderophile elements.<ref>{{cite web|last = Rubin|first = Kenneth H.|url = http://www.soest.hawaii.edu/krubin/GG325/lect34.pdf|title = Lecture 34 Planetary Accretion|publisher=University of Hawaii|access-date = January 16, 2013}}</ref>
{{periodic table (Goldschmidt classification)}}

==Production==

Approximately 100 million [[Tonne|metric tons]] of oxygen are produced yearly. Oxygen is most commonly produced by [[Fractional distillation|fractional distillation]], in which air is cooled to a liquid, then warmed, allowing all the components of air except for oxygen to turn to gases and escape. Fractionally distilling air several times can produce 99.5% pure oxygen.<ref>{{cite encyclopedia|url = https://www.britannica.com/EBchecked/topic/436806/oxygen-O/279404/Commercial-production-and-use|title = Commercial production and use|year = 2013 |access-date=November 25, 2013 |encyclopedia = [[Encyclopædia Britannica]]}}</ref> Another method with which oxygen is produced is to send a stream of dry, clean air through a bed of molecular sieves made of [[Chemistry:Zeolite|zeolite]], which absorbs the nitrogen in the air, leaving 90 to 93% pure oxygen.<ref name = "ReferenceB"/>
[[File:AlbertaSulfurAtVancouverBC.jpg|thumb|Sulfur recovered from oil refining in Alberta, stockpiled for shipment in North Vancouver, British Columbia]]

Sulfur can be mined in its elemental form, although this method is no longer as popular as it used to be. In 1865 a large deposit of elemental sulfur was discovered in the U.S. states of Louisiana and Texas, but it was difficult to extract at the time. In the 1890s, Herman Frasch came up with the solution of liquefying the sulfur with superheated steam and pumping the sulfur up to the surface. These days sulfur is instead more often extracted from [[Chemistry:Oil|oil]], [[Chemistry:Natural gas|natural gas]], and [[Chemistry:Tar|tar]].<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 = http://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}}</ref> Tellurium can also be refined by electrolytic reduction of [[Chemistry:Sodium telluride|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 [[Physics:Isotopes of polonium|isotopes of polonium]] are produced by bombarding [[Chemistry:Bismuth|bismuth]] with neutrons.<ref name="The Elements"/> Polonium can also be produced by high [[Physics:Neutron flux|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"/>

All [[Chemistry:Livermorium|livermorium]] is produced artificially in [[Physics:Particle accelerator|particle accelerator]]s. The first successful production of livermorium was achieved by bombarding curium-248 atoms with [[Physics:Calcium-48|calcium-48]] atoms. As of 2011, roughly 25 atoms of livermorium had been synthesized.<ref name = "ReferenceB"/>

==Applications==

[[Biology:Metabolism|Metabolism]] is the most important source and use of oxygen.  Minor industrial uses include [[Engineering:Steelmaking|Steelmaking]]  (55% of all purified oxygen produced), the [[Engineering:Chemical industry|chemical industry]] (25% of all purified oxygen), medical use, [[Earth:Water treatment|water treatment]] (as oxygen kills some types of bacteria), rocket fuel (in liquid form), and metal cutting.<ref name="ReferenceB"/>

Most sulfur produced is transformed into [[Chemistry:Sulfur dioxide|sulfur dioxide]], which is further transformed into [[Chemistry:Sulfuric acid|sulfuric acid]], a very common industrial chemical. Other common uses include being a key ingredient of [[Chemistry:Gunpowder|gunpowder]] and [[Engineering:Greek fire|Greek fire]], and being used to change [[Biology:Soil pH|soil pH]].<ref name="The Elements"/> Sulfur is also mixed into rubber to vulcanize it. Sulfur is used in some types of [[Chemistry:Concrete|concrete]] and [[Physics:Fireworks|fireworks]]. 60% of all sulfuric acid produced is used to generate [[Chemistry:Phosphoric acid|phosphoric acid]].<ref name="ReferenceB"/><ref>{{cite book|last = le Couteur|first = Penny|title = Napoleon's Buttons|year = 2003|publisher = Penguin Books|isbn=978-1-58542-331-6}}</ref>  Sulfur is used as a [[Biology:Pesticide|pesticide]] (specifically as an [[Chemistry:Acaricide|acaricide]] and [[Biology:Fungicide|fungicide]]) on "orchard, ornamental, vegetable, grain, and other crops."<ref name = EPA-sulfur>{{cite book|title =  Recognition and Management of Pesticide Poisonings|chapter = Other Insecticides and Acaracides|edition = 6th|first1 = James R.|last1 = Roberts|first2 = J. Routt|last2 = Reigart|publisher = Office of Pesticide Programs, U.S. Environmental Protection Agency|year = 2013|page = 93|chapter-url = https://www.epa.gov/sites/production/files/documents/rmpp_6thed_ch9_otherinsecticides.pdf|location = Washington DC}}</ref>
[[File:Black Powder Close Up.jpg|thumb|left|[[Chemistry:Gunpowder|Gunpowder]], an application of sulfur]]

Around 40% of all selenium produced goes to glassmaking. 30% of all selenium produced goes to [[Chemistry:Metallurgy|metallurgy]], including [[Chemistry:Manganese#Occurrence and production|manganese production]]. 15% of all selenium produced goes to [[Earth:Agriculture|agriculture]]. Electronics such as photovoltaic materials claim 10% of all selenium produced. Pigments account for 5% of all selenium produced. Historically, machines such as [[Engineering:Photocopier|photocopier]]s and [[Earth:Light meter|light meter]]s used one-third of all selenium produced, but this application is in steady decline.<ref name="ReferenceB"/>

Tellurium suboxide, a mixture of tellurium and tellurium dioxide, is used in the rewritable data layer of some [[Engineering:CD-RW|CD-RW disks]] and DVD-RW disks. [[Chemistry:Bismuth telluride|Bismuth telluride]] is also used in many microelectronic devices, such as [[Physics:Photodetector|photoreceptors]]. Tellurium is sometimes used as an alternative to sulfur in vulcanized rubber. [[Chemistry:Cadmium telluride|Cadmium telluride]] is used as a high-efficiency material in solar panels.<ref name="ReferenceB"/>

Some of polonium's applications relate to the element's radioactivity. For instance, polonium is used as an [[Physics:Alpha particle|alpha-particle]] generator for research. Polonium alloyed with [[Chemistry:Beryllium|beryllium]] provides an efficient neutron source. Polonium is also used in nuclear batteries. Most polonium is used in antistatic devices.<ref name="ReferenceB"/><ref name="Jackson2002" /> Livermorium does not have any uses whatsoever due to its extreme rarity and short half-life.

Organochalcogen compounds are involved in the [[Physics:Semiconductor|semiconductor]] process. These compounds also feature into [[Chemistry:Ligand|ligand]] chemistry and [[Biology:Biochemistry|biochemistry]]. One application of chalcogens themselves is to manipulate [[Chemistry:Redox|redox]] couples in supramolecular chemistry (chemistry involving non-covalent bond interactions). This application leads on to such applications as crystal packing, assembly of large molecules, and biological recognition of patterns. The secondary bonding interactions of the larger chalcogens, selenium and tellurium, can create organic solvent-holding [[Chemistry:Acetylene|acetylene]] nanotubes. Chalcogen interactions are useful for conformational analysis and stereoelectronic effects, among other things. Chalcogenides with through bonds also have applications. For instance, divalent sulfur can stabilize carbanions, cationic centers, and [[Biology:Radical|radical]]. Chalcogens can confer upon ligands (such as DCTO) properties such as being able to transform [[Chemistry:Copper|Cu(II)]] to Cu(I). Studying chalcogen interactions gives access to radical cations, which are used in mainstream synthetic chemistry. Metallic redox centers of biological importance are tunable by interactions of ligands containing chalcogens, such as [[Chemistry:Methionine|methionine]] and [[Chemistry:Selenocysteine|selenocysteine]]. Also, chalcogen through-bonds{{dubious|date=September 2014}} can provide insight about the process of electron transfer.<ref name = "synth"/>

==Biological role==
{{main|Chemistry:Dioxygen in biological reactions|Biology:Sulfur cycle|Biology:Selenium in biology}}
[[File:DNA chemical structure.svg|thumb|DNA, an important biological compound containing oxygen]]
[[Chemistry:Oxygen|Oxygen]] is needed by almost all organisms for the purpose of generating [[Chemistry:Adenosine triphosphate|ATP]]. It is also a key component of most other biological compounds, such as water, [[Chemistry:Amino acid|amino acid]]s and [[Biology:DNA|DNA]]. Human blood contains a large amount of oxygen. Human bones contain 28% oxygen. Human tissue contains 16% oxygen. A typical 70-kilogram human contains 43 kilograms of oxygen, mostly in the form of water.<ref name="ReferenceB"/>

All animals need significant amounts of [[Chemistry:Sulfur|sulfur]]. Some amino acids, such as [[Chemistry:Cysteine|cysteine]] and [[Chemistry:Methionine|methionine]] contain sulfur. Plant roots take up sulfate ions from the soil and reduce it to sulfide ions. [[Biology:Metalloprotein|Metalloprotein]]s also use sulfur to attach to useful metal atoms in the body and sulfur similarly attaches itself to poisonous metal atoms like [[Chemistry:Cadmium|cadmium]] to haul them to the safety of the liver. On average, humans consume 900 milligrams of sulfur each day. Sulfur compounds, such as those found in skunk spray often have strong odors.<ref name = "ReferenceB"/>

All animals and some plants need trace amounts of [[Chemistry:Selenium|selenium]], but only for some specialized enzymes.<ref name = "The Elements"/><ref name="Selenium: biological information"/> Humans consume on average between 6 and 200 micrograms of selenium per day. Mushrooms and [[Biology:Brazil nut|brazil nut]]s are especially noted for their high selenium content. Selenium in foods is most commonly found in the form of amino acids such as [[Chemistry:Selenocysteine|selenocysteine]] and [[Chemistry:Selenomethionine|selenomethionine]].<ref name = "ReferenceB"/> Selenium can protect against heavy metal poisoning.<ref name = "Selenium: biological information">{{cite web|last = Winter|first= Mark|url = http://www.webelements.com/selenium/biology.html|title =Selenium:Biological information|year = 1993 |access-date=November 25, 2013}}</ref>

Tellurium is not known to be needed for animal life, although a few fungi can incorporate it in compounds in place of selenium. Microorganisms also absorb tellurium and emit [[Chemistry:Dimethyl telluride|dimethyl telluride]]. Most tellurium in the blood stream is excreted slowly in urine, but some is converted to dimethyl telluride and released through the lungs. On average, humans ingest about 600 micrograms of tellurium daily. Plants can take up some tellurium from the soil. Onions and garlic have been found to contain as much as 300 [[Parts per million|parts per million]] of tellurium in dry weight.<ref name="ReferenceB"/>

Polonium has no biological role, and is highly toxic on account of being radioactive.

==Toxicity==
{{NFPA 704|Health = 2|Flammability = 0|Reactivity = 0|S=
|caption=Fire diamond for selenium}}
Oxygen is generally nontoxic, but [[Medicine:Oxygen toxicity|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, sports divers have occasionally drowned from [[Medicine:Convulsion|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 [[Biology:Bacteria|bacteria]].<ref name = "Structure of Matter"/> Ozone, an allotrope of oxygen, is toxic to most life. It can cause [[Medicine:Lesion|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}}</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<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 = http://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 [[Biology:Cattle|cows]] because microbes in the [[Biology:Rumen|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 [[Chemistry:Hydrogen sulfide|hydrogen sulfide]] (H<sub>2</sub>S) and [[Chemistry:Sulfur dioxide|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 [[Biology:Body odor|body odor]]. Extended, low-level exposure, which can occur at some industries, results in weight loss, [[Medicine:Anemia|anemia]], and [[Medicine:Dermatitis|dermatitis]]. In many cases of selenium poisoning, [[Chemistry:Selenous acid|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> [[Chemistry:Hydrogen selenide|Hydrogen selenide]] (H<sub>2</sub>Se) is highly toxic.<ref name="ReferenceB"/>

Exposure to tellurium can produce unpleasant side effects. As little as 10 micrograms of tellurium per cubic meter of air can cause notoriously unpleasant breath, described as smelling like rotten garlic.<ref name="The Elements"/> Acute tellurium poisoning can cause vomiting, gut inflammation, internal bleeding, and respiratory failure. Extended, low-level exposure to tellurium causes tiredness and indigestion. [[Chemistry:Sodium tellurite|Sodium tellurite]] (Na<sub>2</sub>TeO<sub>3</sub>) is lethal in amounts of around 2 grams.<ref name="ReferenceB"/>

Polonium is dangerous as an [[Physics:Alpha particle|alpha particle]] emitter. If ingested, [[Physics:Polonium-210|polonium-210]] is a million times as toxic as [[Chemistry:Hydrogen cyanide|hydrogen cyanide]] by weight; it has been used as a murder weapon in the past, most famously to kill Alexander Litvinenko.<ref name="ReferenceB"/> Polonium poisoning can cause nausea, [[Medicine:Vomiting|vomiting]], anorexia, and lymphopenia. It can also damage hair follicles and [[Biology:White blood cell|white blood cell]]s.<ref name = "ReferenceB"/><ref>{{cite journal|author1=Jefferson, R.D. |author2=Goans R.E. |author3=Blain, P.G. |author4=Thomas, S.H. |pmid=19492929|title = Diagnosis and treatment of polonium poisoning|year = 2009|volume = 47|issue = 5|pages = 379–92|doi = 10.1080/15563650902956431|journal = Clinical Toxicology|s2cid=19648471 }}</ref> Polonium-210 is only dangerous if ingested or inhaled because its alpha particle emissions cannot penetrate human skin.<ref name="Factsheets">{{cite web|url=http://www.iaea.org/Publications/Factsheets/English/polonium210.html |title=Polonium-210 |publisher=International Atomic Energy Agency |year=1998 |access-date=February 11, 2013  |archive-url=https://web.archive.org/web/20120126084112/http://www.iaea.org/Publications/Factsheets/English/polonium210.html |archive-date=January 26, 2012 |df=mdy }}</ref> Polonium-209 is also toxic, and can cause [[Medicine:Leukemia|leukemia]].<ref>{{cite web |last=Freemantle|first=Michael|title=Yasser Arafat, Polonium Poisoning and the Curies|url=http://www.huffingtonpost.co.uk/michael-freemantle/yasser-arafat-polonium-poisoning_b_2199903.html |access-date=November 25, 2013}}</ref>

==Amphid salts==
''Amphid salts'' was a name given by Jons Jacob Berzelius in the 19th century for chemical salts derived from the 16th group of the periodic table which included [[Chemistry:Oxygen|oxygen]], [[Chemistry:Sulfur|sulfur]], [[Chemistry:Selenium|selenium]], and [[Chemistry:Tellurium|tellurium]].<ref>{{Cite book|last=Lee|first=J. D.|url=https://books.google.com/books?id=0m6dIGxmAfwC|title=Concise Inorganic Chemistry, 5TH ed|date=2008-01-03|publisher=Wiley India Pvt. Limited|isbn=978-81-265-1554-7|language=en}}</ref> The term received some use in the early 1800s but is now obsolete.<ref>{{Cite journal|last=Jensen|first=William B.|date=1997-09-01|title=A Note on the Term "Chalcogen"|journal=Journal of Chemical Education|volume=74|issue=9|pages=1063|doi=10.1021/ed074p1063|bibcode=1997JChEd..74.1063J |issn=0021-9584}}</ref> The current term in use for the 16th group is chalcogens.

==See also==
* Chalcogenide
* [[Chemistry:Gold chalcogenides|Gold chalcogenides]]
* [[Chemistry:Halogen|Halogen]]
* [[Chemistry:Interchalcogen|Interchalcogen]]
* [[Chemistry:Pnictogen|Pnictogen]]

==References==
{{Reflist|30em}}

==External links==

{{Periodic table (navbox)}}
{{Navbox periodic table}}
{{Oxides}}
{{Sulfides}}
{{Selenides}}


[[Category:Chalcogens| ]]
[[Category:Periodic table]]
[[Category:Groups (periodic table)]]

{{Sourceattribution|Chalcogen}}