Просмотр кода страницы Английская Википедия:Halogen

{{Английская Википедия/Панель перехода}}
{{short description|Group of chemical elements}}
{{About|the chemical series}}
{{Infobox periodic table group/header
| title       = Halogens
| group number= 17
| trivial name= halogens
| by element  = fluorine group
| CAS         = VIIA
| old IUPAC   = VIIB
| mark        = F,Cl,Br,I,At,Ts
| left        = [[Chalcogen|chalcogens]]
| right       = [[Noble gas|noble gases]]}}
|-
! colspan=2 style="text-align:left;" | ↓&nbsp;[[Period (periodic table)|Period]]
|-
! [[Period 2 element|2]]
| {{element cell image|9|Fluorine|F| |Gas|Halogen|Primordial|image=Liquid fluorine tighter crop.jpg|image caption=Liquid fluorine at cryogenic temperatures}}
|-
! [[Period 3 element|3]]
| {{element cell image|17|Chlorine|Cl| |Gas|Halogen|Primordial|image=Chlorine ampoule.jpg|image caption=Chlorine gas}}
|-
! [[Period 4 element|4]]
| {{element cell image|35|Bromine|Br| |Liquid|Halogen|Primordial|image=Bromine 25ml.jpg|image caption=Liquid bromine}}
|-
! [[Period 5 element|5]]
| {{element cell image|53|Iodine|I| |Solid|Halogen|Primordial|image=Iod kristall.jpg|image caption=Iodine crystal}}
|-
! [[Period 6 element|6]]
| {{element cell image|85|Astatine|At| |Solid|Halogen|From decay}}
|-
! [[Period 7 element|7]]
| {{element cell image|117|Tennessine|Ts| |Unknown phase|Halogen|Synthetic}}
|-
| colspan="2"|
----
''Legend''
{| style="text-align:center; border:0;margin: 0 auto"
|-
| style="border: {{element color|Primordial}}; background:{{Element color|table mark}}; padding:0 2px;" |[[primordial element]]
|-
| style="border: {{element color|from decay}}; background:{{Element color|table mark}};" |[[Trace radioisotope|element from decay]]
|-
| style="border:{{element color|Synthetic}}; background:{{Element color|table mark}};" | [[Synthetic element|Synthetic]]
|-
| Atomic number color:
|-
| <span style="color:{{element color|Solid}}">black=solid</span>, <span style="color:{{element color|Liquid}}">green=liquid</span>, <span style="color:{{element color|Gas}}">red=gas</span>
|}
|}

The '''halogens''' ({{IPAc-en|ˈ|h|æ|l|ə|dʒ|ə|n|,_|ˈ|h|eɪ|-|,_|-|l|oʊ|-|,_|-|ˌ|dʒ|ɛ|n}}<ref>{{cite book |last=Jones |first=Daniel |author-link=Daniel Jones (phonetician) |title=English Pronouncing Dictionary |editor=Peter Roach |editor2=James Hartmann |editor3=Jane Setter |place=Cambridge |publisher=Cambridge University Press |orig-year=1917 |year=2017 |isbn=978-3-12-539683-8 }}</ref><ref>{{MerriamWebsterDictionary|Halogen}}</ref><ref>{{Dictionary.com|Halogen}}</ref>) are a [[group (periodic table)|group]] in the [[periodic table]] consisting of six chemically related [[chemical element|elements]]: [[fluorine]] (F), [[chlorine]] (Cl), [[bromine]] (Br), [[iodine]] (I), and the [[Radioactive decay|radioactive]] elements [[astatine]] (At) and [[tennessine]] (Ts), though some authors<ref>[https://www.researchgate.net/profile/Burkhard-Fricke Fricke, Burkhard] [2007.12.??] [https://www.researchgate.net/publication/225672062_Superheavy_elements_a_prediction_of_their_chemical_and_physical_properties Superheavy elements a prediction of their chemical and physical properties] PDF | "Element 117" | www.researchgate.net | Retrieved - 2023.08.13 (20:58:??) -- yyyy.mm.dd (hh:mm:ss)</ref> would exclude tennessine as its chemistry is unknown and is theoretically expected to be more like that of [[gallium]]. In the modern [[IUPAC]] nomenclature, this group is known as '''group 17'''.<ref name=":0">{{Cite web |title=halogen {{!}} Elements, Examples, Properties, Uses, & Facts {{!}} Britannica |url=https://www.britannica.com/science/halogen |access-date=2022-03-21 |website=www.britannica.com |language=en}}</ref>

The word "halogen" means "salt former" or "salt maker". When halogens react with [[metal]]s, they produce a wide range of [[Salt (chemistry)|salts]], including [[calcium fluoride]], [[sodium chloride]] (common table salt), [[silver bromide]] and [[potassium iodide]].<ref>{{Cite web |title=Chemical properties of the halogens - Group 17 - the halogens - Edexcel - GCSE Combined Science Revision - Edexcel |url=https://www.bbc.co.uk/bitesize/guides/ztjy6yc/revision/2 |access-date=2022-03-21 |website=BBC Bitesize |language=en-GB}}</ref>

The group of halogens is the only [[Group (periodic table)|periodic table group]] that contains elements in three of the main [[State of matter|states of matter]] at [[Standard conditions for temperature and pressure|standard temperature and pressure]], though not far above room temperature the same becomes true of groups [[alkali metals|1]] and [[pnictogen|15]], assuming white phosphorus is taken as the standard state.<ref group="n">This could also be the case for [[Group 12 element|group 12]], although [[copernicium]]'s melting and boiling points are still uncertain.</ref> All of the halogens form acids when bonded to hydrogen. Most halogens are typically produced from [[mineral]]s or [[salt (chemistry)|salts]]. The middle halogens—chlorine, bromine, and iodine—are often used as [[disinfectant]]s. Organobromides are the most important class of [[flame retardant]]s, while elemental halogens are dangerous and can be toxic.
{{toclimit|3}}

== History ==
The fluorine mineral [[fluorospar]] was known as early as 1529. Early chemists realized that fluorine compounds contain an undiscovered element, but were unable to isolate it. In 1860, [[George Gore (chemist)|George Gore]], an English chemist, ran a current of electricity through [[hydrofluoric acid]] and probably produced fluorine, but he was unable to prove his results at the time.{{cn|date=November 2023}} In 1886, [[Henri Moissan]], a chemist in Paris, performed [[electrolysis]] on [[potassium bifluoride]] dissolved in anhydrous [[hydrogen fluoride]], and successfully isolated fluorine.<ref name="Nature's Building Blocks">{{cite book|last=Emsley|first=John|title=Nature's Building Blocks|year=2011|publisher=OUP Oxford |isbn= 978-0199605637}}</ref>

[[Hydrochloric acid]] was known to [[alchemy|alchemists]] and early chemists. However, elemental chlorine was not produced until 1774, when [[Carl Wilhelm Scheele]] heated hydrochloric acid with [[manganese dioxide]]. Scheele called the element "dephlogisticated muriatic acid", which is how chlorine was known for 33 years. In 1807, [[Humphry Davy]] investigated chlorine and discovered that it is an actual element. Chlorine gas was used as a [[poisonous gas]] during [[World War I]]. It displaced oxygen in contaminated areas and replaced common oxygenated air with the toxic chlorine gas. The gas would burn human tissue externally and internally, especially the lungs, making breathing difficult or impossible depending on the level of contamination.<ref name = "Nature's Building Blocks"/>

Bromine was discovered in the 1820s by [[Antoine Jérôme Balard]]. Balard discovered bromine by passing chlorine gas through a sample of [[brine]]. He originally proposed the name ''muride'' for the new element, but the [[French Academy]] changed the element's name to bromine.<ref name = "Nature's Building Blocks"/>

Iodine was discovered by [[Bernard Courtois]], who was using [[seaweed]] ash as part of a process for [[saltpeter]] manufacture. Courtois typically boiled the seaweed ash with water to generate [[potassium chloride]]. However, in 1811, Courtois added sulfuric acid to his process and found that his process produced purple fumes that condensed into black crystals. Suspecting that these crystals were a new element, Courtois sent samples to other chemists for investigation. Iodine was proven to be a new element by [[Joseph Gay-Lussac]].<ref name = "Nature's Building Blocks"/>

In 1931, [[Fred Allison]] claimed to have discovered element 85 with a [[magneto-optical machine]], and named the element Alabamine, but was mistaken. In 1937, [[Rajendralal De]] claimed to have discovered element 85 in minerals, and called the element dakine, but he was also mistaken. An attempt at discovering element 85 in 1939 by [[Horia Hulubei]] and [[Yvette Cauchois]] via [[spectroscopy]] was also unsuccessful, as was an attempt in the same year by [[Walter Minder]], who discovered an iodine-like element resulting from [[beta decay]] of [[polonium]]. Element 85, now named astatine, was produced successfully in 1940 by [[Dale R. Corson]], [[K.R. Mackenzie]], and [[Emilio G. Segrè]], who bombarded [[bismuth]] with [[alpha particle]]s.<ref name = "Nature's Building Blocks"/>

In 2010, a team led by nuclear physicist [[Yuri Oganessian]] involving scientists from the [[JINR]], [[Oak Ridge National Laboratory]], [[Lawrence Livermore National Laboratory]], and [[Vanderbilt University]] successfully bombarded berkelium-249 atoms with calcium-48 atoms to make tennessine.<ref name="117s">{{cite journal|last1=Oganessian |first1=Yu.Ts. |author-link1=Yuri Oganessian |last2=Abdullin |first2=F.Sh. |last3=Bailey |first3=P.D. |last4=Benker |first4=D.E. |last5=Bennett |first5=M.E. |last6=Dmitriev |first6=S.N. |last7=Ezold |first7=J.G. |last8=Hamilton |first8=J.H. |last9=Henderson |first9=R.A. |first10=M.G. |last10=Itkis |first11=Yuri V. |last11=Lobanov |first12=A.N. |last12=Mezentsev |first13=K. J. |last13=Moody |first14=S.L. |last14=Nelson |first15=A.N. |last15=Polyakov | first16=C.E. |last16=Porter |first17=A.V. |last17=Ramayya |first18=F.D. |last18=Riley |first19=J.B. |last19=Roberto |first20=M. A. |last20=Ryabinin | first21=K.P. |last21=Rykaczewski |first22=R.N. |last22=Sagaidak | first23=D.A. |last23=Shaughnessy |first24=I.V. |last24=Shirokovsky |first25=M.A. |last25=Stoyer |first26=V.G. |last26=Subbotin | first27=R. |last27=Sudowe |first28=A.M. |last28=Sukhov |first29=Yu.S. |last29=Tsyganov |first30=Vladimir K. |last30=Utyonkov |first31=A.A. |last31=Voinov |first32=G.K. |last32=Vostokin | first33=P.A. |last33=Wilk |display-authors=6 |title=Synthesis of a new element with atomic number {{nowrap|{{mvar|Z}} {{=}} 117}} |year=2010 |journal=Physical Review Letters |volume=104 |issue=14 |page=142502 |doi=10.1103/PhysRevLett.104.142502 |pmid=20481935 |bibcode=2010PhRvL.104n2502O |s2cid=3263480 |doi-access=free }}</ref>  

=== Etymology ===
In 1811, the German chemist [[Johann Schweigger]] proposed that the name "halogen" – meaning "salt producer", from αλς [hals] "salt" and γενειν [genein] "to beget" – replace the name "chlorine", which had been proposed by the English chemist [[Humphry Davy]].<ref>{{cite journal|last1=Schweigger|first1=J.S.C.|title=Nachschreiben des Herausgebers, die neue Nomenclatur betreffend|journal=Journal für Chemie und Physik|date=1811|volume=3|issue=2|pages=249–255|url=https://babel.hathitrust.org/cgi/pt?id=njp.32101076802287;view=1up;seq=295|trans-title=Postscript of the editor concerning the new nomenclature|language=de}}  On p. 251, Schweigger proposed the word "halogen":  ''"Man sage dafür lieber mit richter Wortbildung ''Halogen'' (da schon in der Mineralogie durch ''Werner's'' Halit-Geschlecht dieses Wort nicht fremd ist) von αλς ''Salz'' und dem alten γενειν (dorisch γενεν) ''zeugen''."'' (One should say instead, with proper morphology, "halogen" (this word is not strange since [it's] already in mineralogy via Werner's "halite" species) from αλς [als] "salt" and the old γενειν [genein] (Doric γενεν) "to beget".)</ref>  Davy's name for the element prevailed.<ref>{{Cite journal|title=J. S. C. Schweigger: His Romanticism and His Crystal Electrical Theory of Matter|journal=Isis|volume=62|issue=3 |pages=328–338|year=1971|doi=10.1086/350763|jstor=229946|last1=Snelders|first1=H. A. M.|s2cid=170337569}}</ref> However, in 1826, the [[Sweden|Swedish]] [[chemist]] Baron [[Jöns Jacob Berzelius]] proposed the term "halogen" for the elements fluorine, chlorine, and iodine, which produce a [[salt (chemistry)|sea-salt]]-like substance when they form a [[chemical compound|compound]] with an alkaline metal.<ref>In 1826, Berzelius coined the terms ''Saltbildare'' (salt-formers) and ''Corpora Halogenia'' (salt-making substances) for the elements chlorine, iodine, and fluorine.  See:  {{cite journal|last1=Berzelius|first1=Jacob|title=Årsberättelser om Framstegen i Physik och Chemie| journal=Arsb. Vetensk. Framsteg |trans-title=Annual Report on Progress in Physics and Chemistry|date=1826|publisher=P.A. Norstedt & Söner|location=Stockholm, Sweden|volume=6|page=187|url=https://babel.hathitrust.org/cgi/pt?id=hvd.32044092556919;view=1up;seq=195|language=sv}} From p. 187: ''"De förre af dessa, d. ä. ''de electronegativa'', dela sig i tre klasser: 1) den första innehåller kroppar, som förenade med de electropositiva, omedelbart frambringa salter, hvilka jag derför kallar ''Saltbildare'' (Corpora Halogenia). Desse utgöras af chlor, iod och fluor *)."'' (The first of them [i.e., elements], the electronegative [ones], are divided into three classes:  1) The first includes substances which, [when] united with electropositive [elements], immediately produce salts, and which I therefore name "salt-formers" (salt-producing substances).  These are chlorine, iodine, and fluorine *).)</ref><ref>The word "halogen" appeared in English as early as 1832 (or earlier).  See, for example: Berzelius, J.J. with A.D. Bache, trans., (1832) [https://books.google.com/books?id=eWsWAAAAYAAJ&pg=RA1-PA248 "An essay on chemical nomenclature, prefixed to the treatise on chemistry,"] ''The American Journal of Science and Arts'', '''22''': 248–276 ; see, for example [https://books.google.com/books?id=eWsWAAAAYAAJ&pg=RA1-PA263 p. 263.]</ref>

The English names of these elements all have the ending [[-ine]]. Fluorine's name comes from the [[Latin]] word ''fluere'', meaning "to flow", because it was derived from the mineral [[fluorite]], which was used as a [[Flux (metallurgy)|flux]] in metalworking. Chlorine's name comes from the [[Greek language|Greek]] word ''chloros'', meaning "greenish-yellow". Bromine's name comes from the Greek word ''bromos'', meaning "stench". Iodine's name comes from the Greek word ''iodes'', meaning "violet". Astatine's name comes from the Greek word ''astatos'', meaning "unstable".<ref name = "Nature's Building Blocks"/> Tennessine is named after the US state of [[Tennessee]].

== Characteristics ==

=== Chemical ===

The halogens fluorine, chlorine, bromine, and iodine are [[Nonmetal (chemistry)|nonmetal]]s; the chemical properties of the two heaviest group&nbsp;17 members have not been conclusively investigated. The halogens show trends in chemical bond energy moving from top to bottom of the periodic table column with fluorine deviating slightly.  It follows a trend in having the highest bond energy in compounds with other atoms, but it has very weak bonds within the diatomic F<sub>2</sub> molecule. This means that further down group&nbsp;17 in the periodic table, the reactivity of elements decreases because of the increasing size of the atoms.<ref>Page 43, Edexcel International GCSE chemistry revision guide, Curtis 2011</ref>

{| class="wikitable" style="float:right; margin-top:0; margin-left:1em; text-align:center; font-size:10pt; line-height:11pt; width:25%;"
|+ style="margin-bottom: 5px;" | Halogen bond energies (kJ/mol){{sfn|Greenwood|Earnshaw|1997|p=804}}
|-
! X
! X<sub>2</sub>
! HX
! BX<sub>3</sub>
! AlX<sub>3</sub>
! CX<sub>4</sub>
|-
! scope="row" style="background:#ff9;"| F
| style="background:#ff9;"| ''159''
| style="background:#ff9;"| ''574''
| style="background:#ff9;"| ''645''
| style="background:#ff9;"| ''582''
| style="background:#ff9;"| ''456''
|-
! scope="row" | Cl
|243
|428
|444
|427
|327
|-
! scope="row" | Br
|193
|363
|368
|360
|272
|-
! scope="row" | I
|151
|294
|272
|285
|239
|}

Halogens are highly [[reactivity (chemistry)|reactive]], and as such can be harmful or lethal to [[Organism|biological organisms]] in sufficient quantities. This high reactivity is due to the high [[electronegativity]] of the atoms due to their high [[effective nuclear charge]]. Because the halogens have seven valence electrons in their outermost energy level, they can gain an electron by reacting with atoms of other elements to satisfy the [[octet rule]]. [[Fluorine]] is the most reactive of all elements; it is the only element more electronegative than oxygen, it attacks otherwise-inert materials such as glass, and it forms compounds with the usually inert [[noble gas]]es. It is a [[Corrosive substance|corrosive]] and highly toxic gas. The reactivity of fluorine is such that, if used or stored in laboratory glassware, it can react with glass in the presence of small amounts of water to form [[silicon tetrafluoride]] (SiF<sub>4</sub>). Thus, fluorine must be handled with substances such as [[polytetrafluoroethylene|Teflon]] (which is itself an [[organofluorine]] compound), extremely dry glass, or metals such as copper or steel, which form a protective layer of fluoride on their surface.

The high reactivity of fluorine allows some of the strongest bonds possible, especially to carbon. For example, Teflon is fluorine bonded with carbon and is extremely resistant to thermal and chemical attacks and has a high melting point.

==== Molecules ====

===== Diatomic halogen molecules =====

The stable halogens form [[homonuclear]] [[diatomic]] [[molecules]].
Due to relatively weak intermolecular forces, chlorine and fluorine form part of the group known as "elemental gases".

{| class="wikitable" style="margin: 1em auto 1em auto; text-align:center;"
|-
! halogen || molecule || structure || model || ''d''(X−X) / pm<br />(gas phase) || ''d''(X−X) / pm<br />(solid phase)
|-
| [[fluorine]] || F<sub>2</sub> || [[Image:Difluorine-2D-dimensions.png|45px]] || [[Image:Fluorine-3D-vdW.png|45px]] || 143 || 149
|-
| [[chlorine]] || Cl<sub>2</sub> || [[Image:Dichlorine-2D-dimensions.png|70px]] || [[Image:Chlorine-3D-vdW.png|63px]] || 199 || 198
|-
| [[bromine]] || Br<sub>2</sub> || [[Image:Dibromine-2D-dimensions.png|80px]] || [[Image:Bromine-3D-vdW.png|72px]] || 228 || 227
|-
| [[iodine]] || I<sub>2</sub> || [[Image:Diiodine-2D-dimensions.png|70px]] || [[Image:Iodine-3D-vdW.png|84px]] || 266 || 272
<!--don't add astatine; it's not proven to form At2 molecules-->
|}

The elements become less reactive and have higher melting points as the atomic number increases. The higher melting points are caused by stronger [[London dispersion force]]s resulting from more electrons.

==== Compounds ====

===== Hydrogen halides =====
{{Main|Hydrogen halides}}
All of the halogens have been observed to react with hydrogen to form [[hydrogen halide]]s. For fluorine, chlorine, and bromine, this reaction is in the form of:

: H<sub>2</sub> + X<sub>2</sub> → 2HX

However, hydrogen iodide and hydrogen astatide can split back into their constituent elements.<ref name = "assorted">{{Cite web|author = Jim Clark|url = http://www.chemguide.co.uk/inorganic/group7/otherreactions.html|title = Assorted reactions of the halogens|year = 2011|access-date = February 27, 2013}}</ref>

The hydrogen-halogen reactions get gradually less reactive toward the heavier halogens. A fluorine-hydrogen reaction is explosive even when it is dark and cold. A chlorine-hydrogen reaction is also explosive, but only in the presence of light and heat. A bromine-hydrogen reaction is even less explosive; it is explosive only when exposed to flames. Iodine and astatine only partially react with hydrogen, forming [[Chemical equilibrium|equilibria]].<ref name = "assorted"/>

All halogens form binary compounds with hydrogen known as the hydrogen halides: [[hydrogen fluoride]] (HF), [[hydrogen chloride]] (HCl), [[hydrogen bromide]] (HBr), [[hydrogen iodide]] (HI), and [[hydrogen astatide]] (HAt). All of these compounds form acids when mixed with water. Hydrogen fluoride is the only hydrogen halide that forms [[hydrogen bond]]s. Hydrochloric acid, hydrobromic acid, hydroiodic acid, and {{not a typo|hydroastatic}} acid are all [[strong acid]]s, but hydrofluoric acid is a [[weak acid]].<ref>{{Cite web|author = Jim Clark|url = http://www.chemguide.co.uk/inorganic/group7/acidityhx.html|title = THE ACIDITY OF THE HYDROGEN HALIDES|year = 2002|access-date = February 24, 2013}}</ref>

All of the hydrogen halides are [[irritation|irritants]]. Hydrogen fluoride and hydrogen chloride are highly [[acid]]ic. Hydrogen fluoride is used as an [[Industry (economics) |indu]][[Manufacturing |strial]] chemical, and is highly toxic, causing [[pulmonary edema]] and damaging cells.<ref>{{Cite web|url = http://www.bt.cdc.gov/agent/hydrofluoricacid/basics/facts.asp|title = Facts about hydrogen fluoride|year = 2005|access-date = 2017-10-28|url-status = dead|archive-url = https://web.archive.org/web/20130201152726/http://www.bt.cdc.gov/agent/hydrofluoricacid/basics/facts.asp|archive-date = 2013-02-01}}</ref> Hydrogen chloride is also a dangerous chemical. Breathing in gas with more than fifty parts per million of hydrogen chloride can cause death in humans.<ref>{{Cite web|url = https://www.cdc.gov/niosh/idlh/7647010.html|title = Hydrogen chloride|access-date = February 24, 2013}}</ref> Hydrogen bromide is even more toxic and irritating than hydrogen chloride. Breathing in gas with more than thirty parts per million of hydrogen bromide can be lethal to humans.<ref>{{Cite web|url = https://www.cdc.gov/niosh/idlh/10035106.html|title = Hydrogen bromide|access-date = February 24, 2013}}</ref> Hydrogen iodide, like other hydrogen halides, is toxic.<ref>{{cite web | url = http://www.kumed.com/~/media/Imported/kumed/documents/kdhehydrogeniodide.ashx | title = Poison Facts:Low Chemicals: Hydrogen Iodid |access-date =2015-04-12}}</ref>

===== Metal halides =====
{{main|Metal halides}}
All the halogens are known to react with sodium to form [[sodium fluoride]], [[sodium chloride]], [[sodium bromide]], [[sodium iodide]], and sodium astatide. Heated sodium's reaction with halogens produces bright-orange flames. Sodium's reaction with chlorine is in the form of:

: {{math|2Na + Cl<sub>2</sub> → 2NaCl}}<ref name = "assorted"/>

Iron reacts with fluorine, chlorine, and bromine to form iron(III) halides. These reactions are in the form of:

: {{math|2Fe + 3X<sub>2</sub> → 2FeX<sub>3</sub>}}<ref name = "assorted"/>

However, when iron reacts with iodine, it forms only [[iron(II) iodide]].

: {{math|Fe + I<sub>2</sub> → FeI<sub>2</sub>}}

Iron wool can react rapidly with fluorine to form the white compound [[iron(III) fluoride]] even in cold temperatures. When chlorine comes into contact with a heated iron, they react to form the black [[iron(III) chloride]]. However, if the reaction conditions are moist, this reaction will instead result in a reddish-brown product. Iron can also react with bromine to form [[iron(III) bromide]]. This compound is reddish-brown in dry conditions. Iron's reaction with bromine is less reactive than its reaction with fluorine or chlorine. A hot iron can also react with iodine, but it forms iron(II) iodide. This compound may be gray, but the reaction is always contaminated with excess iodine, so it is not known for sure. Iron's reaction with iodine is less vigorous than its reaction with the lighter halogens.<ref name = "assorted"/>

===== Interhalogen compounds =====
{{Main|Interhalogen}}

Interhalogen compounds are in the form of XY<sub>n</sub> where X and Y are halogens and n is one, three, five, or seven. Interhalogen compounds contain at most two different halogens. Large interhalogens, such as {{math|ClF<sub>3</sub>}} can be produced by a reaction of a pure halogen with a smaller interhalogen such as {{math|ClF}}. All interhalogens except {{math|[[iodine heptafluoride|IF<sub>7</sub>]]}} can be produced by directly combining pure halogens in various conditions.<ref name = "Chemistry of Interhalogens">{{Cite book|url = https://books.google.com/books?id=nvatWdX1ZWcC|title = Chemistry Of Interhalogen Compounds|year = 2007|access-date = February 27, 2013|isbn = 9788183562430|last1 = Saxena|first1 = P. B| publisher=Discovery Publishing House }}</ref>

Interhalogens are typically more reactive than all diatomic halogen molecules except F<sub>2</sub> because interhalogen bonds are weaker. However, the chemical properties of interhalogens are still roughly the same as those of [[diatomic]] halogens. Many interhalogens consist of one or more atoms of fluorine bonding to a heavier halogen. Chlorine and bromine can bond with up to five fluorine atoms, and iodine can bond with up to seven fluorine atoms. Most interhalogen compounds are [[covalent]] gases. However, some interhalogens are liquids, such as BrF<sub>3</sub>, and many iodine-containing interhalogens are solids.<ref name = "Chemistry of Interhalogens"/>

===== Organohalogen compounds =====
Many synthetic [[organic compounds]] such as [[plastic]] [[polymer]]s, and a few natural ones, contain halogen atoms; these are known as ''halogenated'' compounds or [[organic halide]]s. Chlorine is by far the most abundant of the halogens in seawater, and the only one needed in relatively large amounts (as chloride ions) by humans. For example, chloride ions play a key role in [[brain]] function by mediating the action of the inhibitory transmitter [[Gamma-Aminobutyric acid|GABA]] and are also used by the body to produce stomach acid. Iodine is needed in trace amounts for the production of [[thyroid]] hormones such as [[thyroxine]]. Organohalogens are also synthesized through the [[nucleophilic abstraction]] reaction.<ref name = "Naturally Occurring Organohalogen Compounds-A Comprehensive Update">{{Cite book|url = https://books.google.com/books?id=u45Z-kh61ngC&dq=Organohalogen+compounds&pg=PR1|title = Naturally Occurring Organohalogen Compounds - A Comprehensive Update|year = 2009|access-date = April 23, 2022|isbn = 9783211993224|last1 = Gribble|first1 = G. W| publisher=Springer }}</ref>

===== Polyhalogenated compounds =====

[[Polyhalogenated compound]]s are industrially created compounds substituted with multiple halogens.  Many of them are very toxic and bioaccumulate in humans, and have a very wide application range.  They include [[Polychlorinated biphenyl|PCB]]s, [[PBDE]]s, and [[perfluorinated compound]]s (PFCs), as well as numerous other compounds.

==== Reactions ====

===== Reactions with water =====
Fluorine reacts vigorously with water to produce [[oxygen]] (O<sub>2</sub>) and [[hydrogen fluoride]] (HF):<ref>{{cite web| url = http://www.chemguide.co.uk/inorganic/group7/halogensasoas.html | title = The Oxidising Ability of the Group 7 Elements | publisher =  Chemguide.co.uk|access-date = 2011-12-29}}</ref>

: {{math|2 F<sub>2</sub>(g) + 2 H<sub>2</sub>O(l) → O<sub>2</sub>(g) + 4 HF(aq)}}

Chlorine has maximum solubility of ca. 7.1 g Cl<sub>2</sub> per kg of water at ambient temperature (21&nbsp;°C).<ref>{{cite web| url =http://www.resistoflex.com/chlorine_graphs.htm#9 | title =Solubility of chlorine in water| publisher = Resistoflex.com|access-date =2011-12-29}}</ref>  Dissolved chlorine reacts to form [[hydrochloric acid]] (HCl) and [[hypochlorous acid]], a solution that can be used as a [[disinfectant]] or [[bleach]]:

: {{math|Cl<sub>2</sub>(g) + H<sub>2</sub>O(l) → HCl(aq) + HClO(aq)}}

Bromine has a solubility of 3.41 g per 100 g of water,<ref>{{cite web|url=http://www.bromaid.org/hand_chap1.htm |title=Properties of bromine |publisher=bromaid.org |url-status=dead |archive-url=https://web.archive.org/web/20071208113138/http://www.bromaid.org/hand_chap1.htm |archive-date=December 8, 2007 }}</ref> but it slowly reacts to form [[hydrogen bromide]] (HBr) and [[hypobromous acid]] (HBrO):

: {{math|Br<sub>2</sub>(g) + H<sub>2</sub>O(l) → HBr(aq) + HBrO(aq)}}

Iodine, however, is minimally soluble in water (0.03 g/100 g water at 20&nbsp;°C) and does not react with it.<ref>{{cite web| url =http://hazard.com/msds/mf/baker/baker/files/i2680.htm | title = Iodine MSDS  |publisher=Hazard.com |date =1998-04-21 |access-date =2011-12-29}}</ref>  However, iodine will form an aqueous solution in the presence of iodide ion, such as by addition of [[potassium iodide]] (KI), because the [[triiodide]] ion is formed.

=== Physical and atomic ===
The table below is a summary of the key physical and atomic properties of the halogens. Data marked with question marks are either uncertain or are estimations partially based on [[periodic trends]] rather than observations.

{| class="wikitable" style="text-align:center"
!Halogen
!Standard [[atomic weight]]<br>([[unified atomic mass unit|u]]){{refn|The number given in [[bracket|parentheses]] refers to the [[Uncertainty#In measurements|measurement uncertainty]]. This uncertainty applies to the [[significant figure|least significant figure]](s) of the number prior to the parenthesized value (i.e., counting from rightmost digit to left). For instance, {{val|1.00794|(7)}} stands for {{val|1.00794|0.00007}}, while {{val|1.00794|(72)}} stands for {{val|1.00794|0.00072}}.<ref>{{cite web|url=http://physics.nist.gov/cgi-bin/cuu/Info/Constants/definitions.html|title=Standard Uncertainty and Relative Standard Uncertainty|website=[[CODATA]] reference|publisher=[[National Institute of Standards and Technology]]|access-date=26 September 2011}}</ref>|group=n}}<ref name="atomicweights2009">{{cite journal|last1=Wieser |first1=Michael E.|last2=Coplen|first2=Tyler B. |year=2011|title=Atomic weights of the elements 2009 (IUPAC Technical Report)|journal=[[Pure and Applied Chemistry|Pure Appl. Chem.]] |volume=83|issue=2|pages=359–396|doi=10.1351/PAC-REP-10-09-14|s2cid=95898322|url=http://iupac.org/publications/pac/pdf/2011/pdf/8302x0359.pdf|access-date=5 December 2012 }}</ref>
![[Melting point]]<br>([[kelvin|K]])
![[Melting point]]<br>([[Celsius|°C]])
![[Boiling point]]<br>([[kelvin|K]])<ref name="RubberBible84th">{{cite book | editor = Lide, D. R. | title = CRC Handbook of Chemistry and Physics | edition = 84th | location = Boca Raton, FL |publisher = CRC Press | year = 2003 }}</ref>
![[Boiling point]]<br>([[Celsius|°C]])<ref name="RubberBible84th"/>
![[Density]]<br>(g/cm<sup>3</sup>at 25&nbsp;°C)
![[Electronegativity]]<br>([[Pauling scale|Pauling]])
!First [[ionization energy]]<br>([[Kilojoule per mole|kJ·mol<sup>−1</sup>]]) !! [[Covalent radius]]<br>([[picometre|pm]])<ref name="slater64">{{cite journal|last=Slater|first=J. C.|year=1964|title=Atomic Radii in Crystals|journal=[[Journal of Chemical Physics]]|volume=41|issue=10|pages=3199–3205|bibcode=1964JChPh..41.3199S|doi=10.1063/1.1725697}}</ref>
|-
| [[Fluorine]] || 18.9984032(5) || 53.53 || −219.62 || 85.03 || −188.12 || 0.0017 || 3.98 || 1681.0 || 71
|-
| [[Chlorine]] || [35.446; 35.457]{{refn|The average atomic weight of this element changes depending on the source of the chlorine, and the values in brackets are the upper and lower bounds.<ref name="atomicweights2009"/>|group=n}} || 171.6 || −101.5 || 239.11 || −34.04 || 0.0032 || 3.16 || 1251.2 || 99
|-
| [[Bromine]] || 79.904(1) || 265.8 || −7.3 || 332.0 || 58.8 || 3.1028 || 2.96 || 1139.9 || 114
|-
| [[Iodine]] || 126.90447(3) || 386.85 || 113.7 || 457.4 || 184.3 || 4.933 || 2.66 || 1008.4 || 133
|-
| [[Astatine]] || [210]{{refn|name=unstable|The element does not have any stable [[nuclide]]s, and the value in brackets indicates the [[mass number]] of the longest-lived [[isotope]] of the element.<ref name="atomicweights2009"/>|group=n}} || 575 || 302 || ? 610 || ? 337 || ? 6.2–6.5<ref name=Bonchev>{{cite journal |last1=Bonchev |first1=Danail |last2=Kamenska |first2=Verginia |year=1981 |title=Predicting the properties of the 113–120 transactinide elements |journal=The Journal of Physical Chemistry |volume=85 |issue=9 |pages=1177–86 |doi=10.1021/j150609a021 }}</ref> || 2.2 || 899.0<ref name=1IP>{{cite journal |last1=Rothe |first1=S. |last2=Andreyev |first2=A. N. |last3=Antalic |first3=S. |last4=Borschevsky |first4=A. |last5=Capponi |first5=L. |last6=Cocolios |first6=T. E. |last7=De Witte |first7=H. |last8=Eliav |first8=E. |last9=Fedorov |first9=D. V. | display-authors=8
|year=2013 |title=Measurement of the First Ionization Potential of Astatine by Laser Ionization Spectroscopy |journal=Nature Communications |volume=4 |doi=10.1038/ncomms2819 |url=https://www.researchgate.net/publication/236836716 |last10=Fedosseev |first10=V. N. |last11=Fink |first11=D. A. |last12=Fritzsche |first12=S. |last13=Ghys |first13=L. |last14=Huyse |first14=M. |last15=Imai |first15=N. |last16=Kaldor |first16=U. |last17=Kudryavtsev |first17=Yuri |last18=Köster |first18=U. |last19=Lane |first19=J. F. W. |last20=Lassen |first20=J. |last21=Liberati |first21=V. |last22=Lynch |first22=K. M. |last23=Marsh |first23=B. A. |last24=Nishio |first24=K. |last25=Pauwels |first25=D. |last26=Pershina |first26=V. |last27=Popescu |first27=L. |last28=Procter |first28=T. J. |last29=Radulov |first29=D. |last30=Raeder |first30=S. |pages=1–6 |pmid=23673620 |pmc=3674244|bibcode=2013NatCo...4.1835R }}</ref> || ? 145<ref>{{Cite web|url=https://www.thoughtco.com/astatine-facts-element-ar-606501|title=Get Facts About the Element Astatine|website=www.thoughtco.com|accessdate=November 12, 2021}}</ref>
|-
| [[Tennessine]] || [294]{{refn|name=unstable|group=n}} || ? 623-823<ref name="thoughtco.com">{{Cite web|url=https://www.thoughtco.com/element-117-facts-ununseptium-or-uus-3880071|title=How Much Do You Know About the Element Tennessine?|website=www.thoughtco.com|accessdate=November 12, 2021}}</ref> || ? 350-550<ref name="thoughtco.com"/> || ? 883<ref name="thoughtco.com"/> || ? 610<ref name="thoughtco.com"/> || ? 7.1-7.3<ref name="thoughtco.com"/> || - || ? 743<ref>{{Cite web|url=https://www.webelements.com/tennessine/atoms.html|title=WebElements Periodic Table » Tennessine » properties of free atoms|website=www.webelements.com|accessdate=November 12, 2021}}</ref> || ? 157<ref name="thoughtco.com"/>
|}

{| class="wikitable"
|-
![[Atomic number|Z]] !! [[Chemical element|Element]] !! [[Electron shell|No. of electrons/shell]]
|-
| 9 || fluorine || 2, 7
|-
| 17 || chlorine|| 2, 8, 7
|-
| 35 || bromine || 2, 8, 18, 7
|-
| 53 || iodine || 2, 8, 18, 18, 7
|-
| 85 || astatine || 2, 8, 18, 32, 18, 7
|-
| 117 || tennessine || 2, 8, 18, 32, 32, 18, 7 ''(predicted)''<ref name=Haire>{{Cite book |publisher= [[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>
|}
[[File:Pressure halogens.png|thumb|366x366px|Boiling or sublimation temperature dependence for halogens at various pressures. The vertical bar indicates the melting point]]
{| class="wikitable"
|+Sublimation or boiling point (<sup>o</sup>C) of halogens at various pressures<ref>{{Cite web |title=Краткий справочник физико-химических величин Равделя, Л.: Химия, 1974 г. – 200 стр. \\ стр 67 табл. 24. |url=https://rushim.ru/books/spravochniki/mishenko.pdf}}</ref>
| colspan="2" rowspan="1" |Tmelt (<sup>о</sup>С)
| −100.7
| −7.3
|112.9
|-
!log(P[Pa])
!mmHg
!Cl<sub>2</sub>
!Br<sub>2</sub>
!I<sub>2</sub>
|-
|2.12490302
|1
| −118
| −48.7
|38.7
|-
|2.82387302
|5
| −106.7
| −32.8
|62.2
|-
|3.12490302
|10
| −101.6
| −25
|73.2
|-
|3.42593302
|20
| −93.3
| −16.8
|84.7
|-
|3.72696301
|40
| −84.5
| −8
|97.5
|-
|3.90305427
|60
| −79
| −0.6
|105.4
|-
|4.12490302
|100
| −71.7
|9.3
|116.5
|-
|4.42593302
|200
| −60.2
|24.3
|137.3
|-
|4.72696301
|400
| −47.3
|41
|159.8
|-
|5.00571661
|760
| −33.8
|58.2
|183
|-
!log(P[Pa])
!atm
!Cl<sub>2</sub>
!Br<sub>2</sub>
!I<sub>2</sub>
|-
|5.00571661
|1
| −33.8
|58.2
|183
|-
|5.30674661
|2
| −16.9
|78.8
|
|-
|5.70468662
|5
|10.3
|110.3
|
|-
|6.00571661
|10
|35.6
|139.8
|
|-
|6.30674661
|20
|65
|174
|
|-
|6.48283787
|30
|84.8
|197
|
|-
|6.6077766
|40
|101.6
|215
|
|-
|6.70468662
|50
|115.2
|230
|
|-
|6.78386786
|60
|127.1
|243.5
|
|}

==== Isotopes ====
Fluorine has one stable and naturally occurring [[isotopes of fluorine|isotope]], fluorine-19. However, there are trace amounts in nature of the radioactive isotope fluorine-23, which occurs via [[cluster decay]] of [[protactinium-231]]. A total of eighteen isotopes of fluorine have been discovered, with atomic masses ranging from 13 to 31. 

Chlorine has two stable and naturally occurring [[isotopes of chlorine|isotopes]], chlorine-35 and chlorine-37. However, there are trace amounts in nature of the isotope [[chlorine-36]], which occurs via [[spallation]] of argon-36. A total of 24 isotopes of chlorine have been discovered, with atomic masses ranging from 28 to 51.<ref name = "Nature's Building Blocks"/>

There are two stable and naturally occurring [[isotopes of bromine]], bromine-79 and bromine-81. A total of 33 isotopes of bromine have been discovered, with atomic masses ranging from 66 to 98. 

There is one stable and naturally occurring [[isotopes of iodine|isotope of iodine]], [[iodine-127]]. However, there are trace amounts in nature of the radioactive isotope [[iodine-129]], which occurs via spallation and from the radioactive decay of uranium in ores. Several other radioactive isotopes of iodine have also been created naturally via the decay of uranium. A total of 38 isotopes of iodine have been discovered, with atomic masses ranging from 108 to 145.<ref name = "Nature's Building Blocks"/>

There are no stable [[isotopes of astatine]]. However, there are four naturally occurring radioactive isotopes of astatine produced via radioactive decay of [[uranium]], [[neptunium]], and [[plutonium]]. These isotopes are astatine-215, astatine-217, astatine-218, and astatine-219. A total of 31 isotopes of astatine have been discovered, with atomic masses ranging from 191 to 227.<ref name = "Nature's Building Blocks"/>

There are no stable [[isotopes of tennessine]]. Tennessine has only two known [[synthetic radioisotope]]s, tennessine-293 and tennessine-294.

== Production ==
[[Image:Halogens.jpg|thumb|upright=1.5|right|From left to right: [[chlorine]], [[bromine]], and [[iodine]] at room temperature. Chlorine is a gas, bromine is a liquid, and iodine is a solid. [[Fluorine]] could not be included in the image due to its high [[reactivity (chemistry)|reactivity]], and astatine and tennessine due to their radioactivity.]]
Approximately six million metric tons of the fluorine mineral [[fluorite]] are produced each year. Four hundred-thousand metric tons of hydrofluoric acid are made each year. Fluorine gas is made from hydrofluoric acid produced as a by-product in [[phosphoric acid]] manufacture. Approximately 15,000 metric tons of fluorine gas are made per year.<ref name = "Nature's Building Blocks"/>

The mineral [[halite]] is the mineral that is most commonly mined for chlorine, but the minerals [[carnallite]] and [[sylvite]] are also mined for chlorine. Forty million metric tons of chlorine are produced each year by the [[electrolysis]] of [[brine]].<ref name = "Nature's Building Blocks"/>

Approximately 450,000 metric tons of bromine are produced each year. Fifty percent of all bromine produced is produced in the [[United States]], 35% in [[Israel]], and most of the remainder in [[China]]. Historically, bromine was produced by adding [[sulfuric acid]] and bleaching powder to natural brine. However, in modern times, bromine is produced by electrolysis, a method invented by [[Herbert Dow]]. It is also possible to produce bromine by passing chlorine through seawater and then passing air through the seawater.<ref name = "Nature's Building Blocks"/>

In 2003, 22,000 metric tons of iodine were produced. Chile produces 40% of all iodine produced, [[Japan]] produces 30%, and smaller amounts are produced in [[Russia]] and the United States. Until the 1950s, iodine was extracted from [[kelp]]. However, in modern times, iodine is produced in other ways. One way that iodine is produced is by mixing [[sulfur dioxide]] with [[nitrate]] ores, which contain some [[iodate]]s. Iodine is also extracted from [[natural gas]] fields.<ref name = "Nature's Building Blocks"/>

Even though astatine is naturally occurring, it is usually produced by bombarding bismuth with alpha particles.<ref name = "Nature's Building Blocks"/>

Tennessine is made by using a cyclotron, fusing berkelium-249 and calcium-48 to make tennessine-293 and tennessine-294.

== Applications ==

=== Disinfectants ===
Both chlorine and bromine are used as [[disinfectants]] for drinking water, swimming pools, fresh wounds, spas, dishes, and surfaces. They kill [[bacteria]] and other potentially harmful [[microorganism]]s through a process known as [[Sterilization (microbiology)|sterilization]]. Their reactivity is also put to use in [[Bleaching agent|bleaching]]. [[Sodium hypochlorite]], which is produced from chlorine, is the active ingredient of most [[Cloth|fabric]] bleaches, and chlorine-derived bleaches are used in the production of some [[paper]] products.

=== Lighting ===
[[Halogen lamp]]s are a type of [[incandescent lamp]] using a [[tungsten]] filament in bulbs that have small amounts of a halogen, such as iodine or bromine added. This enables the production of lamps that are much smaller than non-halogen [[incandescent lightbulb]]s at the same [[watt]]age. The gas reduces the thinning of the filament and blackening of the inside of the bulb resulting in a bulb that has a much greater life. Halogen lamps glow at a higher temperature (2800 to 3400 [[kelvin]]) with a whiter colour than other incandescent bulbs. However, this requires bulbs to be manufactured from [[fused quartz]] rather than silica glass to reduce breakage.<ref name="Edison">{{cite web|title=The Halogen Lamp|url=http://www.edisontechcenter.org/halogen.html|website=Edison Tech Center|access-date=2014-09-05 }}</ref>

=== Drug components ===
In [[drug discovery]], the incorporation of halogen atoms into a lead drug candidate results in analogues that are usually more [[lipophilic]] and less water-soluble.<ref>{{cite book| first = G. | last= Thomas | title =Medicinal Chemistry an Introduction | publisher = John Wiley & Sons, West Sussex, UK| year=  2000 | isbn= 978-0-470-02597-0}}</ref> As a consequence, halogen atoms are used to improve penetration through [[lipid membrane]]s and tissues. It follows that there is a tendency for some halogenated drugs to accumulate in [[adipose tissue]].

The chemical reactivity of halogen atoms depends on both their point of attachment to the lead and the nature of the halogen. [[Aromatic]] halogen groups are far less reactive than [[aliphatic]] halogen groups, which can exhibit considerable chemical reactivity. For aliphatic carbon-halogen bonds, the C-F bond is the strongest and usually less chemically reactive than aliphatic C-H bonds. The other aliphatic-halogen bonds are weaker, their reactivity increasing down the periodic table. They are usually more chemically reactive than aliphatic C-H bonds. As a consequence, the most common halogen substitutions are the less reactive aromatic fluorine and chlorine groups.

== Biological role ==
Fluoride anions are found in ivory, bones, teeth, blood, eggs, urine, and hair of organisms. Fluoride anions in very small amounts may be essential for humans.<ref name="who.int">{{cite web|last1=Fawell|first1=J. |title=Fluoride in Drinking-water|url=https://www.who.int/water_sanitation_health/dwq/chemicals/fluoride.pdf|publisher=World Health Organisation|access-date=10 March 2016}}</ref> There are 0.5 milligrams of fluorine per liter of human blood. Human bones contain 0.2 to 1.2% fluorine. Human tissue contains approximately 50 parts per billion of fluorine. A typical 70-kilogram human contains 3 to 6 grams of fluorine.<ref name = "Nature's Building Blocks"/>

Chloride anions are essential to a large number of species, humans included. The concentration of chlorine in the [[dry weight]] of cereals is 10 to 20 parts per million, while in [[potato]]es the concentration of chloride is 0.5%. Plant growth is adversely affected by chloride levels in the [[soil]] falling below 2 parts per million. Human blood contains an average of 0.3% chlorine. Human bone typically contains 900 parts per million of chlorine. Human tissue contains approximately 0.2 to 0.5% chlorine. There is a total of 95 grams of chlorine in a typical 70-kilogram human.<ref name = "Nature's Building Blocks"/>

Some bromine in the form of the bromide anion is present in all organisms. A biological role for bromine in humans has not been proven, but some organisms contain [[organobromine compound]]s. Humans typically consume 1 to 20 milligrams of bromine per day. There are typically 5 parts per million of bromine in human blood, 7 parts per million of bromine in human bones, and 7 parts per million of bromine in human tissue. A typical 70-kilogram human contains 260 milligrams of bromine.<ref name = "Nature's Building Blocks"/>

Humans typically consume less than 100 micrograms of iodine per day. Iodine deficiency can cause [[intellectual disability]]. [[organoiodine chemistry|Organoiodine compound]]s occur in humans in some of the [[gland]]s, especially the [[thyroid gland]], as well as the [[stomach]], [[epidermis (skin)|epidermis]], and [[immune system]]. Foods containing iodine include [[Cod (food)|cod]], [[oyster]]s, [[Shrimp (food)|shrimp]], [[Herring (food)|herring]], [[lobster]]s, [[sunflower seed]]s, [[seaweed]], and [[Edible mushroom|mushrooms]]. However, iodine is not known to have a biological role in plants. There are typically 0.06 milligrams per liter of iodine in human blood, 300 parts per billion of iodine in human bones, and 50 to 700 parts per billion of iodine in human tissue. There are 10 to 20 milligrams of iodine in a typical 70-kilogram human.<ref name = "Nature's Building Blocks"/>

[[Astatine]], although very scarce, has been found in micrograms in the earth.<ref name = "Nature's Building Blocks"/> It has no known biological role because of its high radioactivity, extreme rarity, and has a half-life of just about 8 hours for the most stable isotope. 

Tennessine is purely man-made and has no other roles in nature.

== Toxicity ==
The halogens tend to decrease in toxicity towards the heavier halogens.<ref name = "The Elements"/>

Fluorine gas is extremely toxic; breathing in fluorine at a concentration of 25 parts per million is potentially lethal.  [[Hydrofluoric acid]] is also toxic, being able to penetrate skin and cause [[Hydrofluoric acid burn|highly painful burns]]. In addition, fluoride anions are toxic, but not as toxic as pure fluorine. Fluoride can be lethal in amounts of 5 to 10 grams. Prolonged consumption of fluoride above concentrations of 1.5&nbsp;mg/L is associated with a risk of [[dental fluorosis]], an aesthetic condition of the teeth.<ref>{{cite book|chapter=Guidelines and standards|title=Fluoride in Drinking-water|last1=Fawell | first1=J.|last2= Bailey |first2=K.|last3 =  Chilton |first3= J. |last4=Dahi |first4=E. |last5=Fewtrell |first5=L. |last6=Magara |first6=Y.|publisher=World Health Organization|isbn=978-92-4-156319-2|year=2006|chapter-url=https://www.who.int/water_sanitation_health/publications/fluoride_drinking_water_full.pdf|pages=37–9}}</ref> At concentrations above 4&nbsp;mg/L, there is an increased risk of developing [[skeletal fluorosis]], a condition in which bone fractures become more common due to the hardening of bones. Current recommended levels in [[water fluoridation]], a way to prevent [[dental caries]], range from 0.7 to 1.2&nbsp;mg/L to avoid the detrimental effects of fluoride while at the same time reaping the benefits.<ref>{{cite web|url=https://www.cdc.gov/fluoridation/safety/nrc_report.htm|title=CDC Statement on the 2006 National Research Council (NRC) Report on Fluoride in Drinking Water|date=July 10, 2013|publisher=Centers for Disease Control and Prevention|access-date=August 1, 2013|archive-url=https://web.archive.org/web/20140109183156/http://www.cdc.gov/fluoridation/safety/nrc_report.htm|archive-date=January 9, 2014|url-status=dead|df=mdy-all}}</ref> People with levels between normal levels and those required for skeletal fluorosis tend to have symptoms similar to [[arthritis]].<ref name = "Nature's Building Blocks"/>

Chlorine gas is highly toxic. Breathing in chlorine at a concentration of 3 parts per million can rapidly cause a toxic reaction. Breathing in chlorine at a concentration of 50 parts per million is highly dangerous. Breathing in chlorine at a concentration of 500 parts per million for a few minutes is lethal. In addition, breathing in chlorine gas is highly painful because of its corrosive properties. Hydrochloric acid is the acid of chlorine, while relatively nontoxic, it is highly corrosive and releases very irritating and toxic hydrogen chloride gas in open air.<ref name="The Elements">{{cite book|last=Gray|first=Theodore|title=The Elements|year=2010 |publisher=Running Press | isbn= 9781579128951 }}</ref>

Pure bromine is somewhat toxic but less toxic than fluorine and chlorine. One hundred milligrams of bromine is lethal.<ref name = "Nature's Building Blocks"/> Bromide anions are also toxic, but less so than bromine. Bromide has a lethal dose of 30 grams.<ref name = "Nature's Building Blocks"/>

Iodine is somewhat toxic, being able to irritate the lungs and eyes, with a safety limit of 1 milligram per cubic meter. When taken orally, 3 grams of iodine can be lethal. Iodide anions are mostly nontoxic, but these can also be deadly if ingested in large amounts.<ref name = "Nature's Building Blocks"/>

Astatine is [[radioactive]] and thus highly dangerous, but it has not been produced in macroscopic quantities and hence it is most unlikely that its toxicity will be of much relevance to the average individual.<ref name = "Nature's Building Blocks"/>

Tennessine cannot be chemically investigated due to how short its half-life is, although its radioactivity would make it very dangerous.

== Superhalogen ==
{{Main|Superatom}}
Certain aluminium clusters have superatom properties. These aluminium clusters are generated as anions ({{math|Al{{su|b=''n''|p=−}}}} with ''n'' = 1, 2, 3, ... ) in helium gas and reacted with a gas containing iodine. When analyzed by mass spectrometry one main reaction product turns out to be {{math|{{chem|Al|13|I|−}}}}.<ref name="bergeron/2004">{{cite journal | doi =  10.1126/science.1093902| pmid =  15066775| title =  Formation of Al<sub>13</sub>I<sup>−</sup>: Evidence for the Superhalogen Character of Al<sub>13</sub>| journal =  Science| volume =  304| issue =  5667| pages =  84–7| year =  2004| last1 =  Bergeron| first1 =  D. E.| last2 =  Castleman| first2 =  A. Welford| last3 =  Morisato| first3 =  Tsuguo| last4 =  Khanna| first4 =  Shiv N.| bibcode =  2004Sci...304...84B| s2cid = 26728239}}</ref> These clusters of 13 aluminium atoms with an extra electron added do not appear to react with oxygen when it is introduced in the same gas stream. Assuming each atom liberates its 3 valence electrons, this means 40 electrons are present, which is one of the magic numbers for sodium and implies that these numbers are a reflection of the noble gases.

Calculations show that the additional electron is located in the aluminium cluster at the location directly opposite from the iodine atom. The cluster must therefore have a higher electron affinity for the electron than iodine and therefore the aluminium cluster is called a superhalogen (i.e., the vertical electron detachment energies of the moieties that make up the negative ions are larger than those of any halogen atom).<ref>{{Cite journal|title = Superhalogens as Building Blocks of Halogen-Free Electrolytes in Lithium-Ion Batteries†|journal = Angewandte Chemie|volume = 126|issue = 50|pages = 14136|doi = 10.1002/ange.201408648|year = 2014|last1 = Giri|first1 = Santanab|last2 = Behera|first2 = Swayamprabha|last3 = Jena|first3 = Puru| bibcode=2014AngCh.12614136G }}</ref> The cluster component in the {{math|{{chem|Al|13|I|-}}}} ion is similar to an iodide ion or a bromide ion. The related {{math|{{chem|Al|13|I|2|-}}}} cluster is expected to behave chemically like the [[triiodide]] ion.<ref name="ball">{{cite journal | first =Philip | last = Ball| title =A New Kind of Alchemy| journal =New Scientist | date = 16 April 2005}}</ref><ref name="bergeron/2005">{{cite journal | doi =10.1126/science.1105820| pmid =15653497| title =Al Cluster Superatoms as Halogens in Polyhalides and as Alkaline Earths in Iodide Salts| journal =Science| volume =307| issue =5707| pages =231–5| year =2005| last1 =Bergeron| first1 =D. E.| last2 =Roach| first2 =P. J.| last3 =Castleman| first3 =A. W.| last4 =Jones| first4 =N. O.| last5 =Khanna| first5 =S. N.| bibcode =2005Sci...307..231B| s2cid =8003390}}</ref>

== See also ==
{{wiktionary|halogen}}
* [[Halogen bond]]
* [[Halogen addition reaction]]
* [[Halogen lamp]]
* [[Halogenation]]
* [[Interhalogen]]
* [[Pseudohalogen]]

== Notes ==
{{reflist|group=n}}

== References ==
{{reflist|30em}}

== Further reading ==
* {{Greenwood&Earnshaw2nd}}

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{{Periodic table (navbox)}}
{{Halogens}}
{{fluorides}}
{{chlorides}}
{{bromides}}
{{iodides}}

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[[Category:Halogens| ]]
[[Category:Groups (periodic table)]]
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