Perchloryl fluoride

Perchloryl fluoride
IUPAC name

Perchloryl fluoride
Other names

Chlorine oxyfluoride, Perchlorofluoride, Chlorine fluorine oxide, Trioxychlorofluoride, Perchloric acid fluoride
Identifiers
CAS number 7616-94-6 N
PubChem 24258
ChemSpider 22680 Y
RTECS number SD1925000
Jmol-3D images Image 1
Properties
Molecular formula ClFO3
Molar mass 102.4496 g/mol
Appearance Colorless gas
Density 1.4 g/cm3
Melting point

−147.8 °C, 125 K, -234 °F
Boiling point

−46.7 °C, 226 K, -52 °F
Solubility in water 0.06 g/100 ml (20 °C)
Structure
Molecular shape Tetrahedral[1]:373
Thermochemistry
Std enthalpy of
formation ΔfHo298		 −5.7[1]:380
Hazards
Main hazards Corrosive, oxidizing, toxic
NFPA 704
2
3
3
OX
Threshold Limit Value 3 ppm
 N (verify) (what is: Y/N?)
Except where noted otherwise, data are given for materials in their standard state (at 25 °C, 100 kPa)
Infobox references

Perchloryl fluoride[2] is a reactive gas with the chemical formula FClO3. It has a characteristic sweet odor[3] that resembles gasoline and kerosene. It is toxic and is a powerful oxidizing and fluorinating agent. It is the acid fluoride of perchloric acid.

In spite of its small enthalpy of formation (ΔHfO = −5.7), it is kinetically stable, decomposing only at 400 °C.[1]:380 It is quite reactive towards reducing agents and anions, however, with the chlorine atom acting as an electrophile.[1]:382 It reacts explosively with reducing agents such as amides, metals, hydrides, etc.[3]

Contents

Synthesis and chemistry

Perchloryl fluoride is produced primarily by the fluorination of perchlorates. Antimony pentafluoride is a commonly-used fluorinating agent:[1]:372-373

ClO−
4 + 3 HF + 2 SbF5 → FClO3 + H3O+ + 2 SbF−
6

FClO3 reacts with alcohols to produce alkyl perchlorates, which are extremely shock-sensitive explosives.[4] Using Friedel-Crafts catalysts, it can be used for introducing the –ClO3 group into aromatic rings via electrophilic aromatic substitution.[5]

Applications

Perchloryl fluoride is used in organic chemistry as a mild fluorinating agent.[1]:383 It was the first industrially-relevant electrophilic fluorinating agent, used since the 1960s for producing fluorinated steroids.[4]

Perchloryl fluoride was investigated as a high performance liquid rocket fuel oxidizer.[6] In comparison with chlorine pentafluoride and bromine pentafluoride, it has significantly lower specific impulse, but does not tend to corrode tanks. It does not require cryogenic storage.

It can also be used in flame photometry as an excitation source.[7]

Safety

Perchloryl fluoride is toxic, with a TLV of 3 ppm. It is a strong lung- and eye-irritant capable of producing burns on exposed skin. Its IDLH level is 385 ppm. Symptoms of exposure include dizziness, headaches, syncope, and cyanosis. Exposure to toxic levels causes severe respiratory tract inflammation and pulmonary edema.[6]

References

  1. ^ a b c d e f Harry Julius Emeléus; A. G. Sharpe (1976). Advances in inorganic chemistry and radiochemistry, Volume 18. Academic Press. ISBN 0120236184. 
  2. ^ Chemical Science and Technology Laboratory. "Perchloryl fluoride". National Institute of Standards and Technology. http://webbook.nist.gov/cgi/cbook.cgi?ID=C7616946. Retrieved 2009-11-28. 
  3. ^ a b Jared Ledgard (2007). The Preparatory Manual of Explosives (3rd ed.). Lulu.com. p. 77. ISBN 0615142907. 
  4. ^ a b Peer Kirsch (2004). Modern fluoroorganic chemistry: synthesis, reactivity, applications. Wiley-VCH. p. 74. ISBN 3527306919. 
  5. ^ Peter Bernard David De la Mare (1976). Electrophilic halogenation: reaction pathways involving attack by electrophilic halogens on unsaturated compounds. CUP Archive. p. 63. ISBN 0521290147. 
  6. ^ a b John Burke Sullivan; Gary R. Krieger (2001). Clinical environmental health and toxic exposures (2nd ed.). Lippincott Williams & Wilkins. p. 969. ISBN 068308027X. 
  7. ^ Schmauch, G. E.; Serfass, E. J. (1958). "The Use of Perchloryl Fluoride in Flame Photometry". Applied Spectroscopy 12 (3): 98–102. Bibcode 1958ApSpe..12...98S. doi:10.1366/000370258774615483.  edit