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Concentrated sulfuric acid is one of the most aggressive chemicals in a school laboratory. It chars sugar, dehydrates skin and is used to make hundreds of other chemicals. Yet there’s a family of acids so much stronger that concentrated sulfuric acid looks mild in comparison: acids that can protonate hydrocarbons, turn candle wax into a soup of carbocations, and dissolve substances that most acids can’t touch.
These are the superacids. They’re not just curiosities. They’ve revealed how organic reactions work, won a Nobel Prize, and are used in the petrochemical industry.
The definition
A superacid is any acid stronger than 100% (anhydrous) sulfuric acid.
The definition was proposed by the American chemist James Bryant Conant in 1927, and the field was developed extensively by Ronald Gillespie and George Olah from the 1960s onwards.
Why pH can’t measure them
pH is defined for dilute solutions in water, where pH = −log[H₃O⁺]. That fails for superacids for two reasons:
- Water levels acidity. In water, any acid stronger than H₃O⁺ simply donates its proton to water. So in water, all strong acids look the same. (See the strong acids.)
- Superacids aren’t used in water. They’re used neat, or in non-aqueous solvents, where the idea of [H₃O⁺] doesn’t apply.
Instead, chemists use the Hammett acidity function, H₀, devised by Louis Hammett in 1932. It measures how strongly an acid can protonate a series of very weak indicator bases (such as nitroanilines). In dilute aqueous solutions, H₀ roughly matches pH. For concentrated and non-aqueous acids, it keeps going into negative numbers.
The more negative H₀, the stronger the acid. Each unit is a factor of ten.
| Acid | Approximate H₀ |
|---|---|
| 100% sulfuric acid, H₂SO₄ | −12 |
| Perchloric acid, HClO₄ (anhydrous) | −13 |
| Chlorosulfuric acid, HSO₃Cl | −13.8 |
| Triflic acid, CF₃SO₃H | −14.1 |
| Fluorosulfuric acid, HSO₃F | −15.1 |
| Magic acid (HSO₃F + SbF₅) | about −19 to −23 depending on composition |
| Fluoroantimonic acid (HF + SbF₅) | down to about −28 at high SbF₅ content |
By this measure, fluoroantimonic acid mixtures can be on the order of 10¹⁶ times (ten thousand trillion times) more acidic than pure sulfuric acid, although the exact figure depends on composition and on how it’s measured.
Two types of superacid
1. Brønsted superacids
Single, very strong proton donors. The key feature is an extremely stable conjugate base, so the proton is barely held.
- Triflic acid (trifluoromethanesulfonic acid), CF₃SO₃H. The three fluorine atoms pull electron density away from the sulfonate group, and the negative charge of the triflate ion is spread over three oxygen atoms. Triflate is so stable that triflic acid gives up its proton almost unconditionally. It’s a stable, distillable liquid and one of the most useful superacids in organic synthesis.
- Fluorosulfuric acid, HSO₃F. Similar reasoning: a fluorine in place of one of sulfuric acid’s OH groups.
- Carborane acids, such as H(CHB₁₁Cl₁₁). These are unusual: extremely strong, yet not especially corrosive, because their conjugate base is a large, very stable and unreactive cage of boron atoms. They’ve been described as among the strongest “gentle” acids known.
2. Conjugate (mixed) superacids
A Brønsted acid combined with a strong Lewis acid. The Lewis acid grabs the conjugate base, making it even more stable and freeing the proton even further.
Magic acid is fluorosulfuric acid mixed with antimony pentafluoride (SbF₅). The name comes from a 1966 Christmas party at George Olah’s lab, when a researcher dropped a piece of a candle into it. The wax, which is made of saturated hydrocarbons that normally don’t react with acids at all, dissolved. The mixture had protonated the alkane chains and broken them up into carbocations.
Fluoroantimonic acid is a mixture of hydrogen fluoride and antimony pentafluoride:
2HF + SbF₅ → H₂F⁺ + SbF₆⁻
The SbF₅ takes fluoride from HF to form the very stable hexafluoroantimonate ion (SbF₆⁻), leaving an extremely reactive protonated species, H₂F⁺. The “proton” in this system is effectively free to attack almost anything. See Lewis acids and bases for why SbF₅ is such a powerful electron-pair acceptor.
What superacids can do
Protonate hydrocarbons. Superacids can add a proton to methane:
CH₄ + H⁺ → CH₅⁺
The methanium ion, CH₅⁺, has a carbon atom apparently bonded to five hydrogens, held together by an unusual three-centre, two-electron bond. It’s a real, if fleeting, species studied in superacid media and in the gas phase.
Make carbocations stable enough to study. Carbocations are positively charged carbon species that appear as short-lived intermediates in many organic reactions. In normal conditions they exist for tiny fractions of a second. In superacids, at low temperature, they’re stable enough to be observed directly by NMR spectroscopy. George Olah used this approach to settle long-running debates about carbocation structures, including the famous “non-classical” norbornyl cation. He received the 1994 Nobel Prize in Chemistry for this work.
Protonate very weak bases. Superacids can protonate substances usually considered non-basic, such as carbon monoxide, xenon (in some conditions) and aromatic hydrocarbons.
Industrial uses
- Petroleum refining. Superacid-type catalysts (and solid superacids such as sulfated zirconia) are used to isomerise and crack hydrocarbons, converting straight-chain alkanes into branched alkanes that improve petrol’s octane rating.
- Alkylation. Strong acids including HF and H₂SO₄ are used to combine small alkenes and alkanes into high-octane fuel components.
- Polymer chemistry and pharmaceuticals. Triflic acid and its derivatives are widely used catalysts in organic synthesis.
- Batteries and materials. Triflate and related anions appear in electrolytes because they’re so stable.
Solid superacids
Liquid superacids are difficult to handle and hard to separate from products, so industry increasingly uses solid superacids: materials whose surfaces carry extremely acidic sites. Examples include sulfated zirconia (zirconium dioxide treated with sulfate), certain zeolites, and fluorinated polymer resins such as Nafion, whose sulfonic acid groups sit on a Teflon-like backbone. Reactants flow over the solid, react at the acidic sites and leave, and the catalyst stays in the reactor. This makes processes cleaner and safer, and it’s an active area of green chemistry research.
Handling and hazards
Superacids are among the most dangerous substances in chemistry. Fluoroantimonic acid reacts violently with water, attacks glass (so it must be stored in PTFE containers) and causes catastrophic tissue damage. HF-based superacids also carry the systemic toxicity of hydrogen fluoride. They’re handled only by specialists, in specialised equipment, and never in schools.
Common misconceptions
- “Superacids are just very concentrated acids.” Concentration isn’t the point; they’re fundamentally stronger proton donors, often because a Lewis acid stabilises the conjugate base.
- “The strongest acid must be the most corrosive.” Carborane acids are extremely strong but relatively non-corrosive, because their conjugate base doesn’t react further.
- “Superacids have a pH far below zero.” pH doesn’t apply to them; the Hammett function is used instead.
Key takeaways
- Superacids are acids stronger than 100% sulfuric acid, whose H₀ is about −12.
- Their strength is measured with the Hammett acidity function, not pH.
- Brønsted superacids (like triflic acid) have exceptionally stable conjugate bases; conjugate superacids (like magic acid) add a Lewis acid such as SbF₅.
- They can protonate hydrocarbons and stabilise carbocations, work that earned George Olah the 1994 Nobel Prize.
- For the other extreme, see superbases.
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