Editing Halogen (section)

=== Chemical ===

The halogens fluorine, chlorine, bromine, and iodine are [[Nonmetal (chemistry)|nonmetal]]s; the chemical properties of astatine and tennessine, 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 | archive-date =2012-04-23 | archive-url =https://web.archive.org/web/20120423014139/http://www.resistoflex.com/chlorine_graphs.htm#9 | url-status =dead }}</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.