Why water is not the only stage
By now you have three lenses for acidity — Brønsted proton transfer, Lewis electron-pair donation, and the hard-soft preference picture. Every example so far quietly happened in water. That was convenient, but it hid something important: water is not a neutral referee. It is a chemical participant with opinions, and those opinions blind us to a whole landscape of acid-base behavior. To see that landscape, we have to be willing to pour our reagents into something other than water.
Recall the leveling effect from earlier in this rung. Water can only display acids weaker than its own conjugate acid, H3O+, and bases weaker than its own conjugate base, OH-. Anything stronger than H3O+ simply dumps its proton into water and becomes H3O+; anything more basic than OH- just rips a proton off water and becomes OH-. So in water, HCl, HBr, HNO3, and HClO4 all look identically, boringly strong — flattened to the same ceiling. The full ranking is real; water just refuses to show it. The water acid-base window, from H3O+ down to OH-, is only about 15-16 pKa units wide. Real chemistry runs far past both edges.
There is a second problem: water reacts. Many of the most interesting inorganic species — solvated free electrons, raw carbocations, super-strong bases — would be destroyed the instant they touched water, reducing it, protonating it, or hydrolyzing into mush. So when chemists want to isolate a super-strong acid, dissolve a free electron, or react something that would simply hydrolyze, they deliberately reach for a non-aqueous solvent: liquid ammonia, anhydrous sulfuric acid, liquid HF, glacial acetic acid, molten salts, and others. Each opens a different window and a different stage for reactivity. Choosing the solvent is itself a chemical decision.
The solvent-system concept: a definition for any self-ionizing liquid
How do we even talk about acids and bases once we leave water? One elegant answer generalizes Arrhenius without going all the way to Lewis. Notice that water quietly ionizes itself: 2 H2O gives H3O+ + OH-, producing a characteristic cation and a characteristic anion. The [[solvent-system-concept|solvent-system concept]] says: in ANY solvent that self-ionizes, an acid is whatever raises the concentration of the solvent's own cation, and a base is whatever raises the solvent's own anion. Acid plus base then regenerates neutral solvent — the same neutralization story you already know, just lifted off of water.
Apply this to liquid ammonia, which self-ionizes as 2 NH3 gives NH4+ + NH2- — an ammonium cation and an amide anion, exactly parallel to H3O+ and OH-. By the solvent-system rule, in liquid ammonia an acid is anything that raises NH4+ (so an ammonium salt like NH4Cl is an acid here), and a base is anything that raises NH2- (so sodium amide, NaNH2, is a base). Neutralization in ammonia even looks like its water twin. The pattern keeps working in aprotic liquids: dinitrogen tetroxide ionizes as N2O4 gives NO+ + NO3-, so nitrosyl compounds act as acids and nitrates as bases — and not a proton in sight.
WATER 2 H2O -> H3O+ + OH- acid feeds H3O+ , base feeds OH- AMMONIA 2 NH3 -> NH4+ + NH2- acid feeds NH4+ , base feeds NH2- neutralization in water: HCl + NaOH -> NaCl + H2O neutralization in ammonia: NH4Cl + NaNH2 -> NaCl + 2 NH3 (perfect mirror)
Two solvents at the extremes: liquid ammonia and anhydrous sulfuric acid
Liquid ammonia (which boils at minus 33 degrees C, so it is worked cold) sits on the basic side. It is a far better proton acceptor than water, which means two things. First, it levels acids even harder than water does — almost any acid hands its proton straight to NH3 — so it is no place to rank strong acids. Second, and gloriously, it is a wonderful home for very strong bases. The amide ion NH2- lives happily here. And it has a party trick water can never match: dissolve an alkali metal in it and you get a deep electric-blue solution.
Those famous metal-ammonia solutions are worth a pause. Drop sodium into liquid ammonia and the metal gives up Na+ plus a loose electron — but instead of attacking anything, that electron is caught and cradled by a cage of ammonia molecules, becoming a solvated electron. The blue color is literally the color of dissolved electrons. They are powerful, clean, single-electron reducing agents (chemists use them in the Birch reduction), and they simply cannot exist in water, which the electron would instantly reduce to hydrogen gas. This is the clearest possible demonstration of the chapter's thesis: the solvent decides what can exist.
Now swing to the opposite extreme: anhydrous (100 percent) sulfuric acid. It is so acidic and so unwilling to accept a proton that even nitric acid, a strong acid in water, is forced to behave as a base in it. Sulfuric acid protonates HNO3 and then strips out water: HNO3 + 2 H2SO4 gives NO2+ + H3O+ + 2 HSO4-. The product NO2+ is the nitronium ion, the very electrophile that nitrates benzene in organic chemistry. Read that reaction slowly — the role reversal is the whole point. A substance you have only ever met as an acid is acting as a base, simply because its partner is even more acidic. Acid and base were never fixed labels; they are roles set by the company you keep.
Superacids: stronger than sulfuric acid itself
If pure sulfuric acid is the new reference point, can we beat it? Yes — and the substances that do are called [[superacid|superacids]], defined as any acid stronger than 100 percent sulfuric acid. The ordinary pH scale is useless out here (there is no water to measure), so acidity is reported on the Hammett acidity function H0. Pure sulfuric acid sits near H0 = minus 12; superacids plunge far lower, the strongest reaching around minus 20 and beyond. That number means a proton-donating power that is hard to even picture.
Here is the beautiful trick that makes a superacid, and it ties this whole rung together. You take a strong Brønsted acid and pair it with a strong Lewis acid whose job is to swallow the conjugate base. In 'magic acid' (fluorosulfonic acid plus antimony pentafluoride, HSO3F + SbF5), the SbF5 is a ravenous Lewis acid: it grabs fluoride and SO3F- and locks them into enormous, super-stable, non-nucleophilic anions like SbF6- and Sb2F11-. Those anions are so content that they refuse to hand the proton back. With nothing left to hold it, the proton is stranded, free, and desperate — and acidity skyrockets. The Lewis acid is doing as much of the work as the Brønsted acid.
What does that buy you? Superacids protonate things that are normally not basic at all. Fluoroantimonic acid (HF + SbF5), often cited as the strongest known acid at roughly H0 minus 21 to minus 28, can protonate alkanes — molecules so unreactive that aqueous chemistry treats them as having zero basicity. More than a stunt, this unlocks real science: superacids generate and stabilize carbocations long enough to study them directly (work that won George Olah the 1994 Nobel Prize), and they drive the isomerization and alkylation reactions at the heart of petroleum refining. They are also a vivid mirror of the leveling effect — water could never reveal these acidities, which is exactly why superacid work lives in non-aqueous, usually anhydrous-fluoride media.
Superbases: the mirror image, and what both extremes teach
Acidity has a far edge; so does basicity. A [[superbase|superbase]] is the mirror of a superacid: a base far stronger than the everyday hydroxide ion — strong enough to tear protons off molecules that almost never surrender them, such as the C-H bonds of hydrocarbons. And for exactly the same reason superacids cannot live in water, neither can superbases: water levels any base stronger than OH- (a base that strong simply rips a proton off water to make hydroxide). A trace of moisture destroys them instantly. Superacid and superbase are the two far ends of a single proton-affinity ruler, both walled off by the very same leveling effect.
There is a tidy way to gauge a superbase's strength: a base is as strong as its conjugate acid is weak (recall the reciprocity of conjugate pairs from earlier in this rung). So a base whose conjugate acid is a hopelessly feeble proton-donor must itself be ferocious. Butane has a pKa near 50 — it almost never gives up a proton — so its conjugate base, the butyl anion in n-butyllithium (n-BuLi), is strong enough to deprotonate benzene. The amide ion NH2- (from sodium amide) is a workhorse superbase in liquid ammonia, the hydride ion H- and the oxide ion O2- sit out here too, and chemists prize tailored organic superbases like lithium diisopropylamide (LDA) and the phosphazene 'P4' bases.
One subtlety worth getting right, because it trips people up: basicity is not the same as nucleophilicity. A base wants a proton; a nucleophile wants to bond to a carbon. Bulky superbases like LDA are deliberately engineered to be ferociously strong bases yet poor nucleophiles — their isopropyl groups are too crowded to attack a carbon, so they cleanly pluck a proton without adding to the molecule. That is precisely why a chemist reaches for bulky LDA, rather than small, nucleophilic hydroxide or methoxide, when the goal is to deprotonate a weak carbon acid and make a clean carbanion or enolate.
The big picture: the solvent sets the rules
Step back and the thread is clear. Acid and base are roles, not fixed identities, and which roles a species can play — and how extreme it can get — depends on the stage it stands on. Water is a fine, familiar stage, but it is a narrow one: it levels, it reacts, and its acid-base window is short. Move to liquid ammonia and you can compare strong bases and cradle free electrons; move to anhydrous sulfuric acid and nitric acid becomes a base; add a Lewis acid to mop up the leftover anion and you build a superacid that protonates the unprotonatable; reach for a charge-dense anion kept scrupulously dry and you have a superbase that deprotonates the unprotonatable.
This is also the right moment to puncture a lingering misconception. Stepping outside water does not change a substance's intrinsic nature — it only changes what you can observe and stabilize. HClO4 was always stronger than HCl; water just could not resolve the difference once both passed its H3O+ ceiling. The chemistry was always there; water was simply hiding it. Choosing a non-aqueous medium is how chemists pull back the curtain. With this rung's donor-acceptor lenses in hand, you can now predict not just whether something is an acid or a base, but how far it can be pushed, and in which solvent it must be pushed there.