Fluoroantimonic acid

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Fluoroantimonic acid
H2FSbF6.png
Identifiers
CAS number 16950-06-4 N
ChemSpider 32741664 N
EC number 241-023-8
Jmol-3D images Image 1
Properties
Molecular formula H
2
SbF
7
Molar mass 256.765
Appearance Colorless liquid
Density 2.885 g/cm3
Solubility SO2ClF, SO2
Acidity (pKa) −25
Basicity (pKb) 39
Hazards
GHS hazard statements H300, H310, H314, H330, H411
GHS precautionary statements P260, P264, P273, P280, P284, P301+310
R-phrases R26, R29, R35
S-phrases (S1/2), S36/37/39, S45, S53, S60, S61
Main hazards Extremely corrosive, Violent hydrolysis
NFPA 704
Flammability code 0: Will not burn. E.g., water Health code 4: Very short exposure could cause death or major residual injury. E.g., VX gas Reactivity code 3: Capable of detonation or explosive decomposition but requires a strong initiating source, must be heated under confinement before initiation, reacts explosively with water, or will detonate if severely shocked. E.g., fluorine Special hazard W: Reacts with water in an unusual or dangerous manner. E.g., cesium, sodiumNFPA 704 four-colored diamond
Related compounds
Related acids Antimony pentafluoride

Hydrogen fluoride
Magic acid

Except where noted otherwise, data are given for materials in their standard state (at 25 °C (77 °F), 100 kPa)
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Infobox references

Fluoroantimonic acid (systematically named fluoronium hexafluorostibanide and fluoronium hexafluoroantimonate(V)) is an inorganic compound with the chemical formula H
2
FSbF
6
(also written H
2
F[SbF
6
]
). It is an ionic liquid created by reacting hydrogen fluoride (HF) with antimony pentafluoride (SbF5) in stoichiometrically equivalent amounts. It is the strongest known superacid, which has been demonstrated to protonate even hydrocarbons to afford carbocations and H2.[1] Similar acids can be created by using excess antimony pentafluoride.[2]

The reaction to produce fluoroantimonic acid is:

2 HF \rightleftharpoons H2F+ + F-
SbF5 + F- → SbF6-

The overall reaction is:

SbF5 + 2 HF → SbF6- + H2F+

The second reaction is not in equilibrium, therefore the overall reaction is not in equilibrium. The reaction is exothermic.

The F-, produced by autoionization of hydrogen fluoride, reacts with SbF5 to yield the adduct SbF6-. In the fluoronium ion, hydrogen fluoride is coordinated to the hydrogen ion. The anion is classified as noncoordinating and is both a very weak nucleophile and a very weak base.

The acid is often said to contain "naked protons", but the "free" protons are, in fact, always bonded to hydrogen fluoride molecules to make the fluoronium cations (similar to the hydronium cation in aqueous solution).[3] It is the fluoronium ion that accounts for fluoroantomonic acid's extreme acidity. Fluoroantimonic acid is 1016 (10 quadrillion) times stronger than 100% sulfuric acid.[4] The protons easily migrate through the solution, moving from H2F+ to HF, when present, by the Grotthuss mechanism:

H2F+ + HF \rightleftharpoons HF + H2F+

Fluoroantimonic acid thermally decomposes at higher temperatures, generating hydrogen fluoride gas.

Structure[edit]

Two related products have been crystallised from HF-SbF5 mixtures, and both have been analyzed by single crystal X-ray crystallography. These salts have the formulas [H2F+][Sb2F11] and [H3F2+][Sb2F11]. In both salts, the anion is Sb2F11.[5] As mentioned above, SbF6 is weakly basic; the larger anion Sb2F11 is expected to be still weaker.

The following values[citation needed] are based upon the Hammett acidity function. Increased acidity is indicated by smaller (in this case, more negative) values of H0.

Applications[edit]

This extraordinarily strong acid protonates nearly all organic compounds. In 1967, Bickel and Hogeveen showed that HF-SbF5 will remove H2 from isobutane and methane from neopentane:[6][7]

(CH3)3CH + H+ → (CH3)3C+ + H2
(CH3)4C + H+ → (CH3)3C+ + CH4

Materials compatible with fluoroantimonic acid as a solvent include SO2ClF, and sulfur dioxide; some chlorofluorocarbons have also been used. Containers for HF-SbF5 are made of PTFE.

Safety[edit]

HF-SbF5 has been described as extremely corrosive, toxic, and moisture sensitive.[2]

It reacts violently with water, producing hydrogen fluoride, dioxygen, and antimony (III) fluoride. It will fume in humid air.

See also[edit]

References[edit]

  1. ^ Olah, G. A. (2001). A Life of Magic Chemistry: Autobiographical Reflections of a Nobel Prize Winner. John Wiley and Sons. pp. 100–101. ISBN 0-471-15743-0. 
  2. ^ a b Olah, G. A.; Prakash, G. K. S.; Wang, Q.; Li, X. (2001). "Hydrogen Fluoride–Antimony(V) Fluoride". In Paquette, L. Encyclopedia of Reagents for Organic Synthesis. New York: J. Wiley & Sons. doi:10.1002/047084289X.rh037m. 
  3. ^ Klein, M. L. (October 25, 2000). "Getting the Jump on Superacids" (pdf). Pittsburgh Supercomputing Center (PSC). Retrieved 2012-04-15. 
  4. ^ Olah, G. A. (2005). "Crossing Conventional Boundaries in Half a Century of Research". Journal of Organic Chemistry 70 (7): 2413–2429. doi:10.1021/jo040285o. PMID 15787527. 
  5. ^ Mootz, D.; Bartmann, K. (1988). "The Fluoronium Ions H2F+ and H3F2+: Characterization by Crystal Structure Analysis". Angewandte Chemie, International Edition 27 (3): 391–392. doi:10.1002/anie.198803911. 
  6. ^ Bickel, A. F.; Gaasbeek, C. J.; Hogeveen, H.; Oelderik, J. M.; Platteeuw, J. C. (1967). "Chemistry and spectroscopy in strongly acidic solutions: reversible reaction between aliphatic carbonium ions and hydrogen". Chemical Communications 1967 (13): 634–635. doi:10.1039/C19670000634. 
  7. ^ Hogeveen, H.; Bickel, A. F. (1967). "Chemistry and spectroscopy in strongly acidic solutions: electrophilic substitution at alkane-carbon by protons". Chemical Communications 1967 (13): 635–636. doi:10.1039/C19670000635.