oxygen difluoride
{{chembox
| Watchedfields = changed
| verifiedrevid = 445419652
| Name = Oxygen difluoride
| IUPACName = Oxygen difluoride
| ImageFile = Oxygen-difluoride-2D.png
| ImageSize = 160px
| ImageName = Structure and dimensions of the oxygen difluoride molecule
| ImageFile1 = Oxygen-difluoride-3D-vdW.png
| ImageSize1 = 160px
| ImageName1 = Space-filling model of the oxygen difluoride molecule
| OtherNames = {{Unbulleted list|Oxygen fluoride|Hypofluorous anhydride}}
| Section1 = {{Chembox Identifiers
| ChemSpiderID_Ref = {{chemspidercite|correct|chemspider}}
| ChemSpiderID = 22953
| InChI = 1/F2O/c1-3-2
| InChIKey = UJMWVICAENGCRF-UHFFFAOYAI
| ChEBI_Ref = {{ebicite|correct|EBI}}
| ChEBI = 30494
| SMILES = FOF
| StdInChI_Ref = {{stdinchicite|correct|chemspider}}
| StdInChI = 1S/F2O/c1-3-2
| StdInChIKey_Ref = {{stdinchicite|correct|chemspider}}
| StdInChIKey = UJMWVICAENGCRF-UHFFFAOYSA-N
| CASNo_Ref = {{cascite|correct|??}}
| CASNo = 7783-41-7
| UNII_Ref = {{fdacite|correct|FDA}}
| UNII = 7BCS2CW398
| PubChem = 24547
| RTECS = RS2100000
| EINECS = 231-996-7
}}
| Section2 = {{Chembox Properties
| Formula = {{chem2|OF2}}
| Appearance = colorless gas, pale yellow liquid when condensed
| Odor = peculiar, foul
| MolarMass = 53.9962 g/mol
| Density = {{Unbulleted list|1.90{{nbsp}}g/cm3 (−224{{nbsp}}°C, liquid)|1.719{{nbsp}}g/cm3 (−183{{nbsp}}°C, liquid)|1.521{{nbsp}}g/cm3 (liquid at −145{{nbsp}}°C)|1.88{{nbsp}}g/L (gas at room temperature)}}
| Solvent = slightly soluble in alcohol
| MeltingPtC = −223.8
| BoilingPtC = −144.75
| VaporPressure = 48.9{{nbsp}}atm (at {{convert|−58.0|C|F K|1|disp=or}}{{efn|This is its critical temperature, which is below ordinary room temperature.}})
}}
| Section3 = {{Chembox Structure
| PointGroup = C2V
}}
| Section4 = {{Chembox Thermochemistry
| DeltaHf = 24.5{{nbsp}}kJ mol−1
| DeltaGf = 41.8{{nbsp}}kJ/mol
| Entropy = 247.46{{nbsp}}J/mol K
| HeatCapacity = 43.3{{nbsp}}J/mol K
}}
| Section5 = {{Chembox Hazards
| GHSPictograms = {{GHS03}}{{GHS05}}{{GHS06}}
| GHSSignalWord = danger
| HPhrases = {{HPhrases|H270|H314|H330}}
| PPhrases = {{PPhrases|}}
| GHS_ref = GHS: [https://gestis.dguv.de/data?name=570242 GESTIS 570242]
| NFPA-H = 4
| NFPA-F = 0
| NFPA-R = 3
| NFPA-S = OX
| PEL = TWA 0.05{{nbsp}}ppm (0.1{{nbsp}}mg/m3){{PGCH|0475}}
| LC50 = {{Unbulleted list|2.6{{nbsp}}ppm (rat, 1 hour)|1.5{{nbsp}}ppm (mouse, 1 hour)|26{{nbsp}}ppm (dog, 1 hour)|16{{nbsp}}ppm (monkey, 1 hour)}}{{IDLH|7783417|Oxygen difluoride}}
| REL = C 0.05{{nbsp}}ppm (0.1{{nbsp}}mg/m3)
}}
| Section8 = {{Chembox Related
| OtherCompounds = {{Unbulleted list|Hypofluorous acid|Dioxygen difluoride|difluoramine|Nitrogen trifluoride|Sulfur dichloride|Properties of water|Dichlorine monoxide|Dibromine monoxide|Diiodine monoxide}}}}
}}
Oxygen difluoride is a chemical compound with the formula {{chem2|OF2}}. As predicted by VSEPR theory, the molecule adopts a bent molecular geometry.{{fact|date=December 2024}} It is a strong oxidizer and has attracted attention in rocketry for this reason.{{cite book |doi=10.1016/B0-12-227410-5/00385-9 |chapter=Liquid Rocket Propellants |title=Encyclopedia of Physical Science and Technology |date=2003 |last1=Forbes |first1=Forrest S. |last2=Van Splinter |first2=Peter A. |pages=741–777 |isbn=978-0-12-227410-7 }} With a boiling point of −144.75 °C, OF2 is the most volatile (isolable) triatomic compound.{{Greenwood&Earnshaw2nd|page=819}} The compound is one of many known oxygen fluorides.
Preparation
Oxygen difluoride was first reported in 1929; it was obtained by the electrolysis of molten potassium fluoride and hydrofluoric acid containing small quantities of water.{{cite journal |last1=Lebeau |first1=P. |author-link=Paul Lebeau |last2=Damiens |first2=A. |year=1929 |title=Sur un nouveau mode de préparation du fluorure d'oxygène |trans-title=A new method of preparation of oxygen fluoride |journal=Comptes rendus hebdomadaires des séances de l'Académie des Sciences |volume=188 |pages=1253–1255 |language=fr |access-date=February 21, 2013 |url=http://visualiseur.bnf.fr/CadresFenetre?O=NUMM-3141&I=1253&M=chemindefer}}{{cite journal |last1=Lebeau |first1=P. |author-link=Paul Lebeau |last2=Damiens |first2=A. |year=1927 |title=Sur l'existence d'un composé oxygéné du fluor |trans-title=The existence of an oxygen compound of fluorine |journal=Comptes rendus hebdomadaires des séances de l'Académie des Sciences |volume=185 |pages=652–654 |language=fr |access-date=February 21, 2013 |url=http://visualiseur.bnf.fr/CadresFenetre?O=NUMM-3138&I=652&M=tdm}} The modern preparation entails the reaction of fluorine with a dilute aqueous solution of sodium hydroxide, with sodium fluoride as a side-product:
:{{chem2|2 F2 + 2 NaOH -> OF2 + 2 NaF + H2O}}
Structure and bonding
It is a covalently bonded molecule with a bent molecular geometry and a F-O-F bond angle of 103 degrees. Its powerful oxidizing properties are suggested by the oxidation number of +2 for the oxygen atom instead of its normal −2.
Reactions
Above 200 °C, {{chem2|OF2}} decomposes to oxygen and fluorine by a radical mechanism.
:{{chem2|2 OF2 -> O2 + 2 F2}}
{{chem2|OF2}} reacts with many metals to yield oxides and fluorides. Nonmetals also react: phosphorus reacts with {{chem2|OF2}} to form Phosphorus pentafluoride and {{chem2|POF3}}; sulfur gives sulfur dioxide and sulfur tetrafluoride; and unusually for a noble gas, xenon reacts (at elevated temperatures) yielding Xenon tetrafluoride and xenon oxyfluorides.
Oxygen difluoride reacts with water to form hydrofluoric acid:
:{{chem2|OF2 + H2O -> 2 HF + O2}}
It can oxidize sulfur dioxide to sulfur trioxide and elemental fluorine:
:{{chem2|OF2 + SO2 -> SO3 + F2}}
However, in the presence of UV radiation, the products are sulfuryl fluoride ({{chem2|SO2F2}}) and pyrosulfuryl fluoride ({{chem2|S2O5F2}}):
:{{chem2|OF2 + 2 SO2 -> S2O5F2}}
Safety
{{Expand section|date=August 2018}}
Oxygen difluoride is considered an unsafe gas due to its oxidizing properties. It reacts explosively with water.{{Cite web |title=OXYGEN DIFLUORIDE {{!}} CAMEO Chemicals {{!}} NOAA |url=https://cameochemicals.noaa.gov/chemical/4148 |access-date=2024-05-14 |website=cameochemicals.noaa.gov}} Hydrofluoric acid produced by the hydrolysis of {{chem2|OF2}} with water is highly corrosive and toxic, capable of causing necrosis, leaching calcium from the bones and causing cardiovascular damage, among a host of other highly toxic effects. Other acute poisoning effects include: pulmonary edema, bleeding lungs, headaches, etc.{{Cite web |url=https://www.kdocs.cn/singleSign4CST?cb=https%3A%2F%2Fwww.kdocs.cn%2Fl%2Fcn9j8vXb7Gq3%3Ff%3D201&ts=1715699652 |access-date=2024-05-14 |website=www.kdocs.cn}} Chronic exposure to oxygen difluoride, like that of other chemicals that release fluoride ions, can lead to fluorosis and other symptoms of chronic fluoride poisoning. Oxygen difluoride may be associated with kidney damage. The maximum workplace exposure limit is 0.05 ppm.{{Cite web |title=CDC - NIOSH Pocket Guide to Chemical Hazards - Oxygen difluoride |url=https://www.cdc.gov/niosh/npg/npgd0475.html |access-date=2024-05-14 |website=www.cdc.gov}}
Popular culture
In Robert L. Forward's science fiction novel Camelot 30K, oxygen difluoride was used as a biochemical solvent by fictional life forms living in the solar system's Kuiper belt. While {{chem2|OF2}} would be a solid at 30{{nbsp}}K, the fictional alien lifeforms were described as endothermic, maintaining elevated body temperatures and liquid {{chem2|OF2}} blood by radiothermal heating.
Notes
{{notelist}}
References
{{Reflist|2}}
External links
- [https://web.archive.org/web/20060116134617/http://www.npi.gov.au/database/substance-info/profiles/44.html National Pollutant Inventory - Fluoride and compounds fact sheet]
- [http://webbook.nist.gov/cgi/cbook.cgi?ID=C7783417 WebBook page for {{chem2|OF2}}]
- [https://www.cdc.gov/niosh/npg/npgd0475.html CDC - NIOSH Pocket Guide to Chemical Hazards]
{{Fluorides}}
{{Oxygen compounds}}
{{Fluorine compounds}}