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	{{Chembox		{{Chembox
	| Watchedfields = changed		| Watchedfields = changed
	| verifiedrevid = 443694510		| verifiedrevid = 444658394
			| ImageFile = Difluorine dioxide FOOF.png
	| Watchedfields
	| ImageFileL1 = Fluorine dioxide.svg		| ImageFileL1 = Fluorine dioxide.svg
	| ImageSizeL1 = 120px
	| ImageNameL1 = Stick model of dioxygen difluoride		| ImageNameL1 = Stick model of dioxygen difluoride
	| ImageFileR1 = Dioxygen-difluoride-3D-vdW.png		| ImageFileR1 = Dioxygen-difluoride-3D-vdW.png
	| ImageSizeR1 = 120px
	| ImageNameR1 = Spacefill model of dioxygen difluoride		| ImageNameR1 = Spacefill model of dioxygen difluoride
	| PIN = Dioxygen difluoride		| PIN = Dioxygen difluoride
	| SystematicName = Fluorooxy hypofluorite		| SystematicName = Fluorooxygen hypofluorite
	| OtherNames = Difluorine dioxide<br><!--		| OtherNames = {{plainlist|
			* Monofluorooxygenyl hypofluorite
	-->Fluorine dioxide<br><!--
			* Peroxydifluoride
	-->Perfluoroperoxide
			* Oxygen perfluoride
			* Oxyfluoryl hypofluorite
			* Fluorine peroxide
			* FOOF
			}}
	| Section1 = {{Chembox Identifiers		| Section1 = {{Chembox Identifiers
	| Abbreviations = FOOF		| Abbreviations = FOOF
			| CASNo_Ref = {{cascite|correct|??}}
	| CASNo = 7783-44-0
	| PubChem = 123257		| CASNo = 7783-44-0
			| UNII_Ref = {{fdacite|correct|FDA}}
	| PubChem_Ref = {{Pubchemcite}}
	| ChemSpiderID = 109870		| UNII = XWQ158QF24
			| PubChem = 123257
	| ChemSpiderID_Ref = {{chemspidercite|correct|chemspider}}
			| ChemSpiderID = 109870
	| ChEBI_Ref = {{ebicite|correct|EBI}}
			| ChemSpiderID_Ref = {{chemspidercite|correct|chemspider}}
	| ChEBI = 47866
	| StdInChI_Ref = {{stdinchicite|correct|chemspider}}		| ChEBI_Ref = {{ebicite|correct|EBI}}
	| StdInChI = 1S/F2O2/c1-3-4-2		| ChEBI = 47866
	| StdInChIKey_Ref = {{stdinchicite|correct|chemspider}}		| StdInChI_Ref = {{stdinchicite|correct|chemspider}}
	| StdInChIKey = REAOZOPEJGPVCB-UHFFFAOYSA-N		| StdInChI = 1S/F2O2/c1-3-4-2
			| StdInChIKey_Ref = {{stdinchicite|correct|chemspider}}
	| SMILES = FOOF
	| InChI = 1/F2O2/c1-3-4-2		| StdInChIKey = REAOZOPEJGPVCB-UHFFFAOYSA-N
	| InChIKey = REAOZOPEJGPVCB-UHFFFAOYAK		| SMILES = FOOF
	| Gmelin = 1570		| InChI = 1/F2O2/c1-3-4-2
			| InChIKey = REAOZOPEJGPVCB-UHFFFAOYAK
			| Gmelin = 1570
	}}		}}
	| Section2 =		| Section2 = {{Chembox Properties
			| Formula = {{chem|O|2|F|2}}
	{{Chembox Properties
			| MolarMass = 69.996 g·mol<sup>−1</sup>
	| Formula = {{chem|O|2|F|2}}
			| Appearance = orange as a solid<br />red as a liquid
	| MolarMass = 69.996 g·mol<sup>−1</sup>
	| Solvent = other solvents		| Density = 1.45 g/cm<sup>3</sup> (at b.p.)
			| Solvent = other solvents
	| SolubleOther = decomp.
			| SolubleOther = decomposes
	| MeltingPtC = −154
	| BoilingPtC = −57		| MeltingPtC = −154
			| BoilingPtC = −57
	| Boiling_notes = extrapolated
			| BoilingPt_notes = extrapolated
	}}		}}
	| Section8 =		| Section5 = {{Chembox Thermochemistry
			| DeltaHf = 19.2 kJ/mol
	{{Chembox Related
			| DeltaGf = 58.2 kJ/mol
	| OtherCpds = ]<br><!--
			| Entropy = 277.2 J/(mol·K)
	-->]<br><!--
			| HeatCapacity = 62.1 J/(mol·K)
	-->]<br><!--
			}}
	-->]
			| Section8 = {{Chembox Related
			| OtherCompounds = {{plainlist|
			* {{chem|O|3|F|2|link=Trioxygen difluoride}}
			* {{chem|H|2|O|2|link=Hydrogen peroxide}}
			* {{chem|OF|2|link=Oxygen difluoride}}
			* {{chem|FClO|2|link=Chloryl fluoride}}
			* {{chem|Cl|2|O|2|link=Dichlorine dioxide}}
			* {{chem|S|2|Cl|2|link=Disulfur dichloride}}
			* {{chem|S|2|F|2|link=Disulfur difluoride}}
			}}
			}}
			| Section9 = {{Chembox Hazards
			| NFPA-H = 4
			| NFPA-F = 0
			| NFPA-R = 4
			| NFPA-S = OX
			| GHSPictograms = {{GHS03}}{{GHS05}}{{GHS06}}{{GHS09}}
			| GHSSignalWord = Danger
			| HPhrases = {{H-phrases|}}
			| PPhrases = {{P-phrases|}}
	}}		}}
	}}		}}
			'''Dioxygen difluoride''' is a ] of ] and ] with the ] O<sub>2</sub>F<sub>2</sub>. It can exist as an orange-red colored solid which melts into a red liquid at {{convert|−163|°C|K}}. It is an extremely ] and ] into oxygen and fluorine even at {{convert|−160|°C|K}} at a rate of 4% per {{nowrap|day{{tsp}}{{mdash}}}}{{tsp}}its lifetime at room temperature is thus extremely short.<ref name=Holl>{{cite book| last1=Holleman |first1=A. F.| last2=Wiberg |first2=E.| year=2001| title=Inorganic Chemistry| publisher=]| isbn=978-0-12-352651-9}}</ref> Dioxygen difluoride reacts vigorously with nearly every chemical it encounters (including ordinary ]) leading to its ] nickname '''FOOF''' (a play on its chemical structure and its explosive tendencies).<ref>{{Cite web |last=Lowe |first=Derek |date=2010-02-23 |title=Things I Won't Work With: Dioxygen Difluoride |url=https://www.science.org/content/blog-post/things-i-won-t-work-dioxygen-difluoride |access-date=2022-05-26 |website=www.science.org |language=en}}</ref>
	'''Dioxygen difluoride''' is a ] with the ] {{chem|O|2|F|2}}. It exists as an orange solid that melts into a red liquid at −163&nbsp;&deg;C<ref>
	{{cite journal
	|last=Kirshenbaum |first=A. D.
	|last2=Grosse |first2=A. V.
	|year=1959
	|title=Ozone Fluoride or Trioxygen Difluoride, {{chem|O|3|F|2}}
	|journal=]
	|volume=81 |issue=6 |pages=1277
	|doi=10.1021/ja01515a003
	}}</ref> It is a strong oxidant and decomposes into ] and oxygen even at −160 °C (4% per day).<ref name=Holl>
	{{cite book
	|last=Holleman |first=A. F.
	|last2=Wiberg |first2=E.
	|year=2001
	|title=Inorganic Chemistry
	|publisher=]
	|isbn=0-12-352651-5
	}}</ref>
	==Preparation==		== Preparation ==
	Dioxygen difluoride can be obtained by subjecting a 1:1 mixture of gaseous fluorine and oxygen at low pressure (7–17&nbsp;] is optimal) to an electric discharge of 25–30&nbsp;] at 2.1–2.4&nbsp;]. This is basically the reaction used for the first synthesis by ] in 1933.<ref>		Dioxygen difluoride can be obtained by subjecting a 1:1 mixture of gaseous fluorine and oxygen at low pressure (7–17&nbsp;] (0.9–2.3&nbsp;]) is optimal) to an electric discharge of 25–30&nbsp;] at 2.1–2.4&nbsp;].<ref>{{cite book| first = W. | last = Kwasnik| chapter = Dioxygen Difluoride| title = Handbook of Preparative Inorganic Chemistry| edition = 2nd| editor-first = G. | editor-last = Brauer| publisher = Academic Press| date = 1963| location = NY| volume = 1| page = 162}}</ref>
	{{cite journal
	| last=Ruff | first=O.
	| last2=Mensel | first2=W.
	| year = 1933
	| title = Neue Sauerstofffluoride: {{chem|O|2|F|2}} und OF
	| journal = ]
	| volume = 211 | issue = 1–2 | pages = 204–208
	| doi = 10.1002/zaac.19332110122
	}}</ref> Another synthesis involves mixing O<sub>2</sub> and F<sub>2</sub> in a ] vessel cooled to −196&nbsp;°C, followed by exposing the elements to {{val|3|ul=MeV}} ] for several hours.
			A similar method was used for the first synthesis by ] in 1933.<ref>{{cite journal| last1=Ruff | first1=O.| last2=Mensel | first2=W.| year = 1933
	==Structure and electronic description==
			| title = Neue Sauerstofffluoride: {{chem|O|2|F|2}} und OF| journal = ]| volume = 211 | issue = 1–2 | pages = 204–208| doi = 10.1002/zaac.19332110122}}</ref> Another synthesis involves mixing {{chem|O|2}} and {{chem|F|2}} in a ] vessel cooled to {{convert|-196|°C|K}}, followed by exposing the elements to {{val|3|ul=MeV}} ] for several hours. A third method requires heating a mix of fluorine and oxygen to {{convert|700|C|F}}, and then rapidly cooling it using ].<ref name=thomasmills>{{cite journal |title=Direct synthesis of liquid-phase dioxygen difluoride |first = Thomas | last = Mills | doi = 10.1016/S0022-1139(00)80341-3 | journal = Journal of Fluorine Chemistry | volume = 52 | issue = 3 | year = 1991 | pages = 267–276| url = https://zenodo.org/record/1259619 }}</ref> All of these methods involve synthesis according to the equation
	In {{chem|O|2|F|2}}, oxygen is assigned the unusual ] of +1. In most of its other compounds, oxygen has an oxidation state of −2.
			: {{chem|O|2}} + {{chem|F|2}} → {{chem|O|2|F|2}}
	The structure of dioxygen difluoride resembles that of ], {{chem|H|2|O|2}}, in its large ], which approaches 90°. This geometry conforms with the predictions of ]. The O−O bond length is within 2 pm of the 120.7 ] distance for the O=O double bond in ], O<sub>2</sub>.
			It also arises from the ] of ]:<ref>{{cite journal| last1=Kirshenbaum |first1=A. D.| last2=Grosse |first2=A. V.| year=1959| title=Ozone Fluoride or Trioxygen Difluoride, {{chem|O|3|F|2}}| journal=]| volume=81 |issue=6 |pages=1277| doi=10.1021/ja01515a003}}</ref>
	<center>]</center>
			: 2 {{chem|O|3|F|2}} → 2 {{chem|O|2|F|2}} + {{chem|O|2}}
			== Structure and properties ==
	The bonding within dioxygen difluoride has been the subject of considerable speculation over the years, particularly because of the very short O–O distance and the long O–F distances. Bridgeman has proposed a scheme which essentially has an O–O ''triple'' bond and an O–F single bond that is destabilised and lengthened by repulsion between the ]s on the fluorine atoms and the ]s of the O–O bond.<ref>
			In {{chem|O|2|F|2}}, oxygen is assigned the unusual ] of +1. In most of its other compounds, oxygen has an oxidation state of −2.
	{{cite journal
	|last=Bridgeman |first=A. J.
	|last2=Rothery |first2=J.
	|year=1999
	|title=Bonding in mixed halogen and hydrogen peroxides
	|journal=], ]
	|volume=1999 |issue=22 |pages=4077–4082
	|doi=10.1039/a904968a
	}}</ref> Repulsion involving the fluorine lone pairs is also responsible for the long and weak ].
			The structure of dioxygen difluoride resembles that of ], {{chem|H|2|O|2}}, in its large ], which approaches 90° and C<sub>2</sub> ]. This geometry conforms with the predictions of ].
	==Reactivity==
	The overarching property of this unstable compound is its ] power, despite the fact that all reactions must be conducted near −100 °C.<ref>
			]
	{{cite journal
	|last=Streng |first=A. G.
			The bonding within dioxygen difluoride has been the subject of considerable speculation, particularly because of the very short O−O distance and the long O−F distances. The O−O bond length is within 2 pm of the 120.7 ] distance for the O=O double bond in the ] molecule, {{chem|O|2}}. Several bonding systems have been proposed to explain this, including an O−O ] with O−F single bonds destabilised and lengthened by repulsion between the ]s on the fluorine atoms and the ] of the O−O bond.<ref>{{cite journal| last1=Bridgeman |first1=A. J.
	|year=1963
			| last2=Rothery |first2=J.| year=1999| title=Bonding in mixed halogen and hydrogen peroxides| journal=Journal of the Chemical Society, Dalton Transactions
	|title=The Chemical Properties of Dioxygen Difluoride
			| volume=1999 |issue=22 |pages=4077–4082| doi=10.1039/a904968a}}</ref> Repulsion involving the fluorine lone pairs is also responsible for the long and weak ].
	|journal=]
	|volume=85 |issue=10 |pages=1380–1385
			] indicates that dioxygen difluoride has an exceedingly high barrier to rotation of 81.17&nbsp;kJ/mol around the O−O bond (in hydrogen peroxide the barrier is 29.45&nbsp;kJ/mol); this is close to the O−F bond disassociation energy of 81.59&nbsp;kJ/mol.<ref>{{cite journal |last1=Kraka |first1=Elfi |last2=He |first2=Yuan |last3=Cremer |first3=Dieter |title=Quantum Chemical Descriptions of FOOF: The Unsolved Problem of Predicting Its Equilibrium Geometry |journal=The Journal of Physical Chemistry A |volume=105 |issue=13 |pages=3269–3276 |doi=10.1021/jp002852r |year=2001 |bibcode=2001JPCA..105.3269K}}</ref>
	|doi=10.1021/ja00893a004
	}}</ref> With ] and ], it gives the corresponding ] salts:<ref name=Holl/><ref>{{cite journal
			The ] ] of dioxygen difluoride is 865&nbsp;ppm, which is by far the highest chemical shift recorded for a fluorine nucleus, thus underlining the extraordinary electronic properties of this compound. Despite its instability, thermochemical data for {{Chem|O|2|F|2}} have been compiled.<ref>{{Cite journal |title = Thermodynamic Properties of Dioxygen Difluoride (O<sub>2</sub>F<sub>2</sub>) and Dioxygen Fluoride (O<sub>2</sub>F) |url = https://www.nist.gov/data/PDFfiles/jpcrd364.pdf |last = Lyman |first = John L. |publisher = American Chemical Society and the American Institute of Physics for the National Institute of Standards and Technology |year = 1989 |access-date = 5 August 2013 |journal = Journal of Physical and Chemical Reference Data|volume = 18 |issue = 2 |page = 799 |doi = 10.1063/1.555830 |bibcode = 1989JPCRD..18..799L |archive-date = 4 March 2016 |archive-url = https://web.archive.org/web/20160304080316/http://www.nist.gov/data/PDFfiles/jpcrd364.pdf |url-status = dead }}</ref>
	|last=Solomon |first=I. J.
	|coauthors=''et al.''
			== Reactivity ==
	|year=1964
			The compound readily ] into oxygen and fluorine. Even at a temperature of {{convert|−160|°C|K}}, 4% decomposes each day<ref name=Holl /> by this process:
	|title=New Dioxygenyl Compounds
	|journal=]
			: {{chem|O|2|F|2}} → {{chem|O|2}} + {{chem|F|2}}
	|volume=3 |issue=3 |pages=457
	|doi=10.1021/ic50013a036
			The other main property of this unstable compound is its ] power, although most experimental reactions have been conducted near {{convert|−100|°C|K}}.<ref name=streng>{{cite journal| last=Streng |first=A. G.| year=1963| title=The Chemical Properties of Dioxygen Difluoride| journal=]| volume=85 |issue=10 |pages=1380–1385| doi=10.1021/ja00893a004}}</ref> Several experiments with the compound resulted in a series of fires and explosions. Some of the compounds that produced violent reactions with {{chem|O|2|F|2}} include ], ], ], and even ].<ref name="streng" />
	}}</ref>
			With {{chem|BF|3|link=boron trifluoride}} and {{chem|PF|5|link=phosphorus pentafluoride}}, it gives the corresponding ] salts:<ref name="Holl" /><ref>{{cite journal| last1 = Solomon | first1 = Irvine J.| first2 = Robert I. | last2 = Brabets| first3 = Roy K. | last3 = Uenishi| first4 = James N. | last4 = Keith| first5 = John M. | last5 = McDonough| year=1964| title=New Dioxygenyl Compounds| journal=]
			| volume=3 |issue=3 |pages=457| doi=10.1021/ic50013a036}}</ref>
	:2 {{chem|O|2|F|2}} + 2 {{chem|PF|5}} → 2 {{chem||+||-}} + {{chem|F|2}}		:2 {{chem|O|2|F|2}} + 2 {{chem|PF|5}} → 2 {{chem||+||-}} + {{chem|F|2}}
			== Uses ==
	It converts uranium and plutonium oxides into the corresponding hexafluorides.<ref>
			The compound currently has no practical applications, but has been of theoretical interest. ] used it to synthesize ] at unprecedentedly low temperatures, which was significant because previous methods for preparation needed temperatures so high that the plutonium hexafluoride created would decompose rapidly.<ref>{{cite journal| last1 = Malm | first1 = J. G.| first2 = P. G. | last2 = Eller| first3 = L. B. | last3 = Asprey| title = Low temperature synthesis of plutonium hexafluoride using dioxygen difluoride| journal = ]| volume = 106| issue = 9| date = 1984| pages = 2726–2727| doi=10.1021/ja00321a056}}</ref>
	{{cite book
	|last=Atwood |first=D. A.
	|year=2006
	|chapter=Fluorine: Inorganic Chemistry
	|title=]
	|publisher=]
	|doi=10.1002/0470862106.ia076
	|isbn=
	}}</ref>
	==References==		== See also ==
			* ]
	{{reflist}}
			* ]
	==External links==		== References ==
			{{reflist|30em}}
	*
	*
	*
			== External links ==
	]
			* {{nist | id = C7783440 | name = Perfluoroperoxide }}
	]
			{{Oxygen compounds}}
			{{Fluorine compounds}}
			{{fluorides}}
			]
			]
	]		]
	]		]
			]
			]
	]
	]
	]
	]
	]
	]