Fluoroantimonic acid
| Fluoroantimonic acid | |
|---|---|
 | |
| IUPAC name Fluoroantimonic acid | |
| Identifiers | |
| CAS number | 16950-06-4  |
| PubChem | 6337100 |
| ChemSpider | 21241496 |
| EC number | 241-023-8 |
| Jmol-3D images | {{#if:F[Sb-](F)(F)(F)(F)[FH+]|Image 1 |
| |
| |
| Properties | |
| Molecular formula | H 2SbF 7 |
| Molar mass | 236.76 g mol-1 |
| 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 |
 0
4
3
|
| Related compounds | |
| Related acids | Antimony pentafluoride |
(verify) (what is:  /?)Except where noted otherwise, data are given for materials in their standard state (at 25 °C (77 °F), 100 kPa) | |
| Infobox references | |
Fluoroantimonic acid (systematically named fluorium hexafluorostibanuide and fluorium hexafluoridoantimonate(1-)) is an inorganic compound with the chemical formula H
2FSbF
6 (also written H
2F[SbF
6]). It is an ionic liquid created by reacting hydrogen fluoride with antimony pentafluoride in stoichiometrically equivalent amounts. Similar acids can be created by using excess antimony pentafluoride.[1] The 1:1 combination forms the strongest known superacid, which has been demonstrated to protonate even hydrocarbons to afford carbocations and H2.[2]
The reaction of hydrogen fluoride (HF) and SbF5 is exothermic. F-, being a homoassociation product Lewis base of hydrogen fluoride, attacks the molecules of SbF5 to give an adduct. In the fluoronium molecule, the hydrogen fluoride is coordinated to the hydrogen, although the molecule is formally classified as noncoordinating, because it is both a very weak nucleophile and a very weak base.
Despite the proton being called effectively "naked," it is in fact always attached to a hydrogen fluoride molecule through a very weak dative bond, similar to the hydronium cation.[3] However, the weakness of this bond accounts for the system's extreme acidity. Fluoroantimonic acid is 1016 (10 quadrillion) times stronger than 100% sulfuric acid.[4] The acidic proton easily jumps among different anion-clusters, for example, by the Grotthuss mechanism:
- H2F+ + HF
HF + H2F+
The reaction to produce fluoroantimonic acid is:
- 2 HF H2F+ + F-
- SbF5 + F- → SbF6-
Overall:
- SbF5 + 2 HF → SbF6- + H2F+
Note that the second reaction is not in equilibrium, hence, the overall reaction is not in equilibrium. Fluoroantimonic acid thermally decomposes at higher temperatures, however, emitting hydrogen fluoride vapour.
Structure
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 classified as weakly basic; the larger monoanion Sb2F11− would be expected to be still weaker.
Comparison with other acids
The following values[citation needed] are based upon the Hammett acidity function. Acidity is indicated by large negative values of H0.
- Fluoroantimonic acid (1990) (H0 Value = −31.3)
- Magic acid (1974) (H0 Value = −19.2)
- Carborane superacid (1969) (H0 Value = −18.0)
- Fluorosulfuric acid (1944) (H0 Value = −15.1)
- Triflic acid (1940) (H0 Value = −14.9)
Applications
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
In terms of materials compatible with fluoroantimonic acid as solvents include SO2ClF, and sulfur dioxide; some chlorofluorocarbons have also been used. Containers for HF-SbF5 are made of PTFE.
Safety
HF-SbF5 has been described as extremely corrosive, toxic, and moisture sensitive.[1]
It reacts violently with water, producing hydrogen fluoride, dioxygen, and trifluoridoantimony. As such, it will fume in humid air.
See also
References
- ↑ 1.0 1.1 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.
- ↑ 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.
- ↑ Klein, M. L. (October 25, 2000). "Getting the Jump on Superacids" (pdf). Pittsburgh Supercomputing Center (PSC). Retrieved 2012-04-15.
- ↑ 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.
- ↑ 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.
- ↑ 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.
- ↑ 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.