From Arrhenius to a proton's-eye view
The first picture of acids you ever met was almost certainly the [[arrhenius-acid-base|Arrhenius]] one: an acid releases H+ in water, a base releases OH-. It is true as far as it goes, but it is shackled to water and it leaves out far too much. Ammonia, NH3, is a base in everyone's intuition, yet it has no OH- to give. To set inorganic acidity free we widen the lens, and the [[inorg-bronsted-lowry-acid-base|Brønsted-Lowry]] definition does exactly that with a single, almost austere idea: an acid is a proton donor and a base is a proton acceptor. That is all. Forget the solvent, forget oxygen — focus only on the bare proton, H+, changing hands.
A proton is special precisely because it is so naked: strip the single electron off a hydrogen atom and nothing is left but a point of positive charge, smaller than any other cation by a factor of roughly a hundred thousand. It cannot float around alone in solution; it clamps onto whatever lone pair it can find. So in water the "H+" we casually write is really the hydronium ion, H3O+, a proton riding on a water molecule's lone pair. Keeping that in mind makes the whole Brønsted machinery click: every acid-base event is just a lone pair on the base reaching out and capturing the proton the acid lets go.
Conjugate pairs: every acid hides a base
Here is the move that makes the Brønsted picture so powerful. When an acid donates its proton, what is left behind is itself a base — because that leftover species can, in principle, grab the proton back. The acid and its de-protonated remnant form a [[inorg-conjugate-acid-base-pair|conjugate acid-base pair]], two species differing by exactly one H+. HCl gives up a proton to become Cl-, so HCl and Cl- are a conjugate pair; HCl is the acid, Cl- is its conjugate base. Acetic acid CH3COOH pairs with acetate CH3COO-. And because water can play either role, it appears in two pairs at once: as an acid it sheds H+ to give OH-, and as a base it accepts H+ to become H3O+.
Watch this on a single line and the symmetry jumps out. Write "HCl + H2O gives Cl- + H3O+": HCl (acid) and Cl- (its conjugate base) are one pair, while H2O (base) and H3O+ (its conjugate acid) are the other — every proton transfer always produces two conjugate pairs at once. Now write "NH3 + H2O gives NH4+ + OH-": here NH3 is the base and water is the acid, with NH4+/NH3 and H2O/OH- as the two pairs. The same H2O is the acid in the second equation and the base in the first, which is the textbook example of an amphoteric solvent.
The pairing carries a quantitative payoff: the stronger an acid, the weaker its conjugate base, and vice versa, in a strict reciprocal relationship. A strong acid like HCl hands its proton off almost completely, so Cl- has essentially no appetite to take it back — Cl- is a vanishingly weak base. Flip it around and a strong base, like the amide ion NH2-, has a very weak conjugate acid (NH3). This reciprocity is not a vague slogan; it is exact, because the same proton-transfer equilibrium is being read forwards for the acid and backwards for its conjugate base. Learn to see acid and base as two ends of one seesaw and half of acid-base chemistry becomes obvious by inspection.
The leveling effect: why water hides the strongest acids
Ask which is the stronger acid, HCl or HBr or HClO4, and in water you cannot tell — they all read as "strong," fully dissociated. That is not because they are equally strong; it is because the solvent itself sets a ceiling. Any acid stronger than H3O+ simply dumps its proton onto water completely, and the most acidic species that can actually exist in water is therefore H3O+ itself. This ceiling is the [[inorg-leveling-effect|leveling effect]]: water levels every super-strong acid down to the strength of hydronium, the way a low doorway makes everyone taller than it stoop to the same height. There is a matching floor at the bottom — any base stronger than OH- just rips a proton off water and is leveled to hydroxide.
So how do chemists ever rank the strong acids against each other? They change the solvent. Move to glacial acetic acid, a much weaker proton acceptor than water, and the strong acids are no longer all leveled — they spread out, because acetic acid is reluctant enough that only the very strongest acids fully protonate it. This is the whole logic behind a [[superacid|superacid]], a medium more acidic than pure sulfuric acid: you must escape water entirely, because water would mask the very strength you are trying to create. The leveling effect is a beautiful warning that "strong" and "weak" are not absolute labels stamped on a molecule; they are statements about an acid relative to whatever base the solvent provides.
Oxoacids and Pauling's rules: strength written in structure
Now to the inorganic acids that fill the laboratory shelf: the [[oxoacid|oxoacids]]. These have the general shape (HO)mEOn — a central atom E carrying some hydroxyl OH groups (whose H atoms are the acidic protons) and some terminal, doubly-bonded oxygens with no hydrogen attached. Sulfuric acid is (HO)2SO2; nitric acid is (HO)NO2; phosphoric acid is (HO)3PO. The acidic proton is never bonded straight to the central atom — it always hangs off an oxygen, and the question of acid strength is simply: how willingly does that O-H bond let its proton go?
Linus Pauling spotted that the answer is governed almost entirely by n, the count of terminal oxygens, giving the celebrated [[paulings-rules-for-acidity|Pauling's rules]]. Rule one: for the first ionization, the acid's pKa is roughly 8 - 5n. With no terminal oxygens (n = 0, as in HOCl) pKa is about 8 — a weak acid. One terminal oxygen (n = 1, as in HNO2 or H3PO4) drops pKa to roughly 3, moderately strong. Two (n = 2, as in H2SO4) gives pKa near -3, strong; three (n = 3, as in HClO4) drives it to about -8, a textbook strong acid. Rule two: each successive proton is about 5 pKa units harder to remove than the last, because pulling a second H+ off an already-negative anion is uphill work.
Pauling rule 1: pKa1 ~= 8 - 5n (n = number of terminal =O) n=0 HOCl pKa ~ +7.5 very weak n=1 HNO2,H3PO4 pKa ~ +2 to 3 weak/moderate n=2 H2SO4 pKa ~ -3 strong n=3 HClO4 pKa ~ -8 very strong Pauling rule 2: each further proton pKa increases by ~5 H3PO4 pKa1~2.1 pKa2~7.2 pKa3~12.4
Why should counting oxygens work so well? Because acid strength is really about stabilizing the conjugate base. Each terminal oxygen is an electron-hungry sink that pulls electron density away from the central atom and, after ionization, helps spread the new negative charge over several oxygens by resonance. The more terminal oxygens, the more the anion's charge is delocalized, the more stable the conjugate base, and the more freely the proton leaves. Electronegativity of the central atom matters too — for the same n, a more electronegative E gives a stronger acid (HOCl beats HOBr beats HOI) — but the terminal-oxygen count is the dominant dial, which is why one integer captures so much.
Reading strength off a formula — and where the rules bend
Pauling's rules turn a formula into a strength estimate by a quick procedure, but the trap is that you must read the real structure, not the way the formula happens to be written.
- Draw the oxoacid in (HO)mEOn form: put every acidic H on its own oxygen, then see how many oxygens are left as terminal (=O) with no hydrogen.
- Count the terminal oxygens to get n, and estimate the first pKa from pKa1 ~= 8 - 5n.
- For each additional acidic proton, add roughly 5 to the previous pKa.
- Compare acids with the same n by central-atom electronegativity — the more electronegative E wins.
One last honest framing before you climb on. The Brønsted picture is itself just one lens — the lens of the moving proton. It is superb for the oxoacids and for anything in solution, but it has nothing to say about a reaction with no proton in sight, such as BF3 grabbing the lone pair of NH3. That broader donor-acceptor view is the [[inorg-lewis-acid-base|Lewis]] picture, the very next guide in this rung, and it will swallow the Brønsted definition as a special case. For now, hold onto the core idea: acidity is the willingness to give up a proton, conjugate pairs tie every acid to a base, the solvent sets the ruler you measure on, and for an oxoacid you can read the answer almost straight off its structure.