In chemistry, a superacid (according to the original definition) is an acid with an acidity greater than that of 100% pure sulfuric acid (H2SO4),[1] which has a Hammett acidity function (H0) of −12. According to the modern definition, a superacid is a medium in which the chemical potential of the proton is higher than in pure sulfuric acid.[2] Commercially available superacids include trifluoromethanesulfonic acid (CF3SO3H), also known as triflic acid, and fluorosulfuric acid (HSO3F), both of which are about a thousand times stronger (i.e. have more negative H0 values) than sulfuric acid. Most strong superacids are prepared by the combination of a strong Lewis acid and a strong Brønsted acid. A strong superacid of this kind is fluoroantimonic acid. Another group of superacids, the carborane acid group, contains some of the strongest known acids. Finally, when treated with anhydrous acid, zeolites (microporous aluminosilicate minerals) will contain superacidic sites within their pores. These materials are used on massive scale by the petrochemical industry in the upgrading of hydrocarbons to make fuels.[citation needed]
History
editThe term superacid was originally coined by James Bryant Conant in 1927 to describe acids that were stronger than conventional mineral acids.[1] This definition was refined by Ronald Gillespie in 1971, as any acid with an H0 value lower than that of 100% sulfuric acid (−11.93).[3] George A. Olah prepared the so-called "magic acid", so named for its ability to attack hydrocarbons, by mixing antimony pentafluoride (SbF5) and fluorosulfonic acid (FSO3H).[4] The name was coined after a candle was placed in a sample of magic acid after a Christmas party. The candle dissolved, showing the ability of the acid to protonate alkanes, which under normal acidic conditions do not protonate to any extent.
At 140 °C (284 °F), FSO3H–SbF5 protonates methane to give the tertiary-butyl carbocation, a reaction that begins with the protonation of methane:[4]
- CH4 H → CH
5 - CH
5 → CH
3 H2 - CH
3 3 CH4 → (CH3)3C 3H2
Common uses of superacids include providing an environment to create, maintain, and characterize carbocations. Carbocations are intermediates in numerous useful reactions such as those forming plastics and in the production of high-octane gasoline.
Origin of extreme acid strength
editTraditionally, superacids are made from mixing a Brønsted acid with a Lewis acid. The function of the Lewis acid is to bind to and stabilize the anion that is formed upon dissociation of the Brønsted acid, thereby removing a proton acceptor from the solution and strengthening the proton donating ability of the solution. For example, fluoroantimonic acid, nominally (H
2FSbF
6), can produce solutions with a H0 lower than –28, giving it a protonating ability over a billion times greater than 100% sulfuric acid.[5][6] Fluoroantimonic acid is made by dissolving antimony pentafluoride (SbF5) in anhydrous hydrogen fluoride (HF). In this mixture, HF releases its proton (H ) concomitant with the binding of F− by the antimony pentafluoride. The resulting anion (SbF−
6) delocalizes charge effectively and holds onto its electron pairs tightly, making it an extremely poor nucleophile and base. The mixture owes its extraordinary acidity to the weakness of proton acceptors (and electron pair donors) (Brønsted or Lewis bases) in solution. Because of this, the protons in fluoroantimonic acid and other superacids are popularly described as "naked", being readily donated to substances not normally regarded as proton acceptors, like the C–H bonds of hydrocarbons. However, even for superacidic solutions, protons in the condensed phase are far from being unbound. For instance, in fluoroantimonic acid, they are bound to one or more molecules of hydrogen fluoride. Though hydrogen fluoride is normally regarded as an exceptionally weak proton acceptor (though a somewhat better one than the SbF6– anion), dissociation of its protonated form, the fluoronium ion H2F to HF and the truly naked H is still a highly endothermic process (ΔG° = 113 kcal/mol), and imagining the proton in the condensed phase as being "naked" or "unbound", like charged particles in a plasma, is highly inaccurate and misleading.[7]
More recently, carborane acids have been prepared as single component superacids that owe their strength to the extraordinary stability of the carboranate anion, a family of anions stabilized by three-dimensional aromaticity, as well as by electron-withdrawing group typically attached thereto.
In superacids, the proton is shuttled rapidly from proton acceptor to proton acceptor by tunneling through a hydrogen bond via the Grotthuss mechanism, just as in other hydrogen-bonded networks, like water or ammonia.[8]
Applications
editIn petrochemistry, superacidic media are used as catalysts, especially for alkylations. Typical catalysts are sulfated oxides of titanium and zirconium or specially treated alumina or zeolites. The solid acids are used for alkylating benzene with ethene and propene as well as difficult acylations, e.g. of chlorobenzene.[9] In organic chemistry, superacids are used as a means of protonating alkanes to promote the use of carbocations in situ during reactions. The resulting carbocations are of much use in organic synthesis of numerous organic compounds, the high acidity of the superacids helps to stabilize the highly reactive and unstable carbocations for future reactions.
Examples
editThe following are examples of superacids. Each is listed with its Hammett acidity function,[10] where a smaller value of H0 (in these cases, more negative) indicates a stronger acid.
- Fluoroantimonic acid (HF:SbF5, H0 = -28)
- Magic acid (HSO3F:SbF5, H0 = −23)
- Carborane acids (H(HCB11X11), H0 ≤ −18, indirectly determined and depends on substituents)
- Fluoroboric acid (HF:BF3, H0 = −16.6)
- Bistriflimidic acid (NH(CF3SO2)2, H0 = -15.8.[11] Estimated value calculated from pKa values in 1,2-dichloroethane in comparison to triflic acid)
- Fluorosulfuric acid (FSO3H, H0 = −15.1)
- Hydrogen fluoride (HF, H0 = −15.1)[12]
- Triflic acid (HOSO2CF3, H0 = −14.9)
- Oleum (SO3:H2SO4, H0 = −14.5)[13]
- Perchloric acid (HClO4, H0 = −13)
- Sulfuric acid (H2SO4, H0 = −11.9)
See also
editReferences
edit- ^ a b Hall NF, Conant JB (1927). "A Study of Superacid Solutions". Journal of the American Chemical Society. 49 (12): 3062–70. doi:10.1021/ja01411a010.
- ^ Himmel D, Goll SK, Leito I, Krossing I (2010). "A Unified pH Scale for All Phases". Angew. Chem. Int. Ed. 49 (38): 6885–6888. doi:10.1002/anie.201000252. PMID 20715223.
- ^ Gillespie, R. J.; Peel, T. E.; Robinson, E. A. (1971-10-01). "Hammett acidity function for some super acid systems. I. Systems H2SO4-SO3, H2SO4-HSO3F, H2SO4-HSO3Cl, and H2SO4-HB(HSO4)4". Journal of the American Chemical Society. 93 (20): 5083–5087. doi:10.1021/ja00749a021. ISSN 0002-7863.
The work of Jorgenson and Hartter formed the basis for the present work, the object of which was to extend the range of acidity function measurements into the super acid region, i.e., into the region of acidities greater than that of 100% H2SO4.
- ^ a b George A. Olah, Schlosberg RH (1968). "Chemistry in Super Acids. I. Hydrogen Exchange and Polycondensation of Methane and Alkanes in FSO3H–SbF5 ("Magic Acid") Solution. Protonation of Alkanes and the Intermediacy of CH5 and Related Hydrocarbon Ions. The High Chemical Reactivity of "Paraffins" in Ionic Solution Reactions". Journal of the American Chemical Society. 90 (10): 2726–7. doi:10.1021/ja01012a066.5+ and Related Hydrocarbon Ions. The High Chemical Reactivity of "Paraffins" in Ionic Solution Reactions&rft.volume=90&rft.issue=10&rft.pages=2726-7&rft.date=1968&rft_id=info:doi/10.1021/ja01012a066&rft.au=George A. Olah, Schlosberg RH&rfr_id=info:sid/en.wikipedia.org:Superacid" class="Z3988">
- ^ Olah, George 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.
- ^ Herlem, Michel (1977). "Are reactions in superacid media due to protons or to powerful oxidising species such as SO3 or SbF5?". Pure and Applied Chemistry. 49: 107–113. doi:10.1351/pac197749010107. S2CID 98483167.
- ^ Ruff, F. (Ferenc) (1994). Organic reactions : equilibria, kinetics, and mechanism. Csizmadia, I. G. Amsterdam: Elsevier. ISBN 0444881743. OCLC 29913262.
- ^ Schneider, Michael (2000). "Getting the Jump on Superacids". Pittsburgh Supercomputing Center. Archived from the original on 23 August 2018. Retrieved 20 November 2017.
- ^ Michael Röper, Eugen Gehrer, Thomas Narbeshuber, Wolfgang Siegel "Acylation and Alkylation" in Ullmann's Encyclopedia of Industrial Chemistry, Wiley-VCH, Weinheim, 2000. doi:10.1002/14356007.a01_185
- ^ Gillespie, R. J.; Peel, T. E. (1973-08-01). "Hammett acidity function for some superacid systems. II. Systems sulfuric acid-[fsa], potassium fluorosulfate-[fsa], [fsa]-sulfur trioxide, [fsa]-arsenic pentafluoride, [sfa]-antimony pentafluoride and [fsa]-antimony pentafluoride-sulfur trioxide". Journal of the American Chemical Society. 95 (16): 5173–5178. doi:10.1021/ja00797a013. ISSN 0002-7863.
- ^ Fuller, Maurice (2022). Coordination Chemistry and its Application (PDF). Bibliotex. pp. 45, 46.
- ^ Liang, Joan-Nan Jack (1976). The Hammett Acidity Function for Hydrofluoric Acid and some related Superacid Systems (Ph.D. Thesis, advisor: R. J. Gillespie) (PDF). Hamilton, Ontario: McMaster University. p. 109.
- ^ Olah, George (2009). Superacid Chemistry (PDF). John Wiley & Sons, Inc. p. 47.