Inchemistry, asuperacid (according to the original definition) is anacid with an acidity greater than that of 100% puresulfuric acid (H2SO4),[1] which has aHammett acidity function (H0) of −12. According to the modern definition, a superacid is a medium in which thechemical potential of theproton is higher than in pure sulfuric acid.[2] Commercially available superacids includetrifluoromethanesulfonic acid (CF3SO3H), also known as triflic acid, andfluorosulfuric acid (HSO3F), both of which are about a thousand times stronger (i.e. have more negativeH0 values) than sulfuric acid. Most strong superacids are prepared by the combination of a strongLewis acid and a strongBrønsted acid. A strong superacid of this kind isfluoroantimonic acid. Another group of superacids, thecarborane acid group, contains some of the strongest known acids. Finally, when treated withanhydrous acid,zeolites (microporous aluminosilicate minerals) will contain superacidic sites within their pores.[3] These materials are used on massive scale by the petrochemical industry in the upgrading of hydrocarbons to make fuels.[4]
The termsuperacid was originally coined byJames Bryant Conant in 1927 to describe acids that were stronger than conventionalmineral acids.[1] This definition was refined byRonald Gillespie in 1971, as any acid with anH0 value lower than that of 100% sulfuric acid (−11.93).[5]George A. Olah prepared the so-called "magic acid", so named for its ability to attackhydrocarbons, by mixingantimony pentafluoride (SbF5) andfluorosulfonic acid (FSO3H).[6] 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 toprotonatealkanes, which under normal acidic conditions do not protonate to any extent.
At 140 °C (284 °F), FSO3H–SbF5 protonatesmethane to give the tertiary-butylcarbocation, a reaction that begins with the protonation of methane:[6]
Common uses of superacids include providing an environment to create, maintain, and characterizecarbocations. Carbocations are intermediates in numerous useful reactions such as those forming plastics and in the production ofhigh-octanegasoline.
Traditionally, 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 aH0 lower than −28, giving it a protonating ability over a billion times greater than 100% sulfuric acid.[7][8] Fluoroantimonic acid is made by dissolvingantimony pentafluoride (SbF5) in anhydroushydrogen fluoride (HF). In this mixture, HF releases its proton (H+) concomitant with the binding of F− by the antimony pentafluoride. The resultinganion (SbF−
6) delocalizes charge effectively and holds onto its electron pairs tightly, making it an extremely poornucleophile andbase. 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, theprotons 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 theSbF−
6 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.[9]
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 theGrotthuss mechanism, just as in other hydrogen-bonded networks, like water or ammonia.[10]
Inpetrochemistry, superacidic media are used as catalysts, especially foralkylations. Typical catalysts are sulfated oxides oftitanium andzirconium or specially treated alumina orzeolites. Thesolid acids are used for alkylating benzene withethene andpropene as well as difficultacylations, e.g. ofchlorobenzene.[11] Inorganic chemistry, superacids are used as a means of protonating alkanes to promote the use of carbocationsin 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.
The following are examples of superacids. Each is listed with itsHammett acidity function,[12] where a smaller value ofH0 (in these cases, more negative) indicates a stronger acid.
{{cite book}}: CS1 maint: multiple names: editors list (link)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.