Acid-base concepts

Ask a school student what makes something an acid and you will usually hear "it releases $H^+$ in water." That answer is not wrong — it is just the first of three widening circles chemists drew around the idea of "acid" and "base" over about eighty years. Svante Arrhenius drew the first, narrow circle in 1884: acids and bases only exist in water, defined by $H^+$ and $OH^-$ . Johannes Brønsted and Thomas Lowry widened it in 1923 to any proton transfer, in any solvent or none at all. Gilbert Lewis, that same year, drew the widest circle of all: an acid–base reaction is simply one species offering a pair of electrons and another accepting it — no proton required. Each theory does not overturn the one before it; it swallows it whole and then covers more ground. Every Arrhenius acid is a Brønsted–Lowry acid, and every Brønsted–Lowry acid is a Lewis acid, but the reverse is not true. This article builds the three definitions in that order, shows how to spot conjugate acid–base pairs, and works through why $\text{BF}_3 + \text{NH}_3$ is acid–base chemistry even though no water and no $H^+$ are anywhere in sight.

The core insight

Arrhenius, Brønsted–Lowry, and Lewis are three nested acid–base definitions — proton-in-water, proton donor/acceptor, and electron-pair acceptor/donor — each strictly broader than the last, so every reaction classified by an earlier theory is automatically classified the same way by a later one, but not vice versa.

The core idea

Arrhenius (1884) — the narrowest circle. An acid raises $[H^+]$ (equivalently $[H_3O^+]$ ) when dissolved in water; a base raises $[OH^-]$ in water. This is intuitive and matches everyday examples like $\text{HCl}$ and $\text{NaOH}$ , but it is tied to one solvent. It cannot explain why gaseous $\text{NH}_3$ , which contains no $OH^-$ group at all, behaves as a base.

Brønsted–Lowry (1923) — proton transfer, any solvent. An acid is a proton donor; a base is a proton acceptor. Every such reaction is a reversible proton hand-off:

\text{acid}_1 + \text{base}_2 \rightleftharpoons \text{base}_1 + \text{acid}_2

where: acid $_1$ and base $_1$ are a conjugate pair (they differ by exactly one $H^+$ ), and base $_2$ /acid $_2$ are the other conjugate pair. The species that donates the proton becomes the conjugate base; the species that accepts it becomes the conjugate acid. This single idea explains why water is amphiprotic — it can donate a proton to a stronger base or accept one from a stronger acid, playing acid in one reaction and base in the next.

Lewis (1923) — electron pairs, no proton needed. A Lewis acid is an electron-pair acceptor (it needs an empty or low-energy orbital); a Lewis base is an electron-pair donor (it needs a lone pair). This is the same donor/acceptor language used for coordinate (dative) bonds, and it is deliberately broader: it covers boron and aluminium compounds, metal cations, and any reaction that forms a new coordinate bond, whether or not a proton moves at all. Because $H^+$ itself has an empty orbital and accepts an electron pair from a base, every Brønsted–Lowry acid–base reaction is automatically a valid Lewis acid–base reaction too — the nesting is exact, not approximate.

Figure

Three nested circles labelled Lewis outermost then Bronsted Lowry then Arrhenius innermost with example species placed in the ring where each first qualifies as acid base chemistry

Why it works

Why does each new theory strictly contain the old one instead of just being a different, incompatible way of talking about acids? The answer is that each theory generalises the mechanism of the one before it while keeping every reaction the earlier theory already explained. Arrhenius chemistry is really proton transfer to water: $\text{HCl} + \text{H}_2\text{O} \rightleftharpoons \text{H}_3\text{O}^+ + \text{Cl}^-$ is nothing but an acid donating a proton to the base water. Brønsted and Lowry noticed that nothing in that mechanism requires the acceptor to be water — any proton acceptor works, in any solvent, or with no solvent at all (e.g., gaseous $\text{HCl} + \text{NH}_3 \rightarrow \text{NH}_4\text{Cl}$ ). So every Arrhenius acid–base pair is a special case of Brønsted–Lowry proton transfer where the base happens to be water. Lewis then asked what a proton transfer actually is at the electronic level: the base's lone pair forms a new bond to the proton's empty $1s$ orbital. That is exactly the definition of a Lewis acid–base reaction, with $H^+$ playing the role of the (extremely small) Lewis acid. Once you see a proton transfer as "lone pair moves into an empty orbital," it is obvious that the same mechanism works when the empty orbital belongs to boron, aluminium, or a metal cation instead of a bare proton — and the Lewis definition captures exactly that extra ground, with no proton involved at all. This is also why the general strength rule holds across all three theories: the stronger an acid, the weaker its conjugate base (a strong proton donor leaves behind a species with little pull for reclaiming that proton), because donating and accepting are just opposite descriptions of the same electron-pair/proton event.

Worked example

Classify a reaction two ways: identify the Brønsted acid, base, and both conjugate pairs, then confirm the same reaction also satisfies the Lewis definition. Consider

\text{NH}_3 + \text{H}_2\text{O} \rightleftharpoons \text{NH}_4^+ + \text{OH}^-

Step 1. Decide which species loses a proton and which gains one, reading left to right. Water converts to $\text{OH}^-$ , which means it has lost an $H^+$ — so $\text{H}_2\text{O}$ is the Brønsted acid here. Ammonia converts to $\text{NH}_4^+$ , gaining an $H^+$ — so $\text{NH}_3$ is the Brønsted base.

Step 2. Pair each species with the partner it became by gaining or losing exactly one proton. $\text{H}_2\text{O}$ (acid) and $\text{OH}^-$ (its conjugate base) are one conjugate pair; $\text{NH}_3$ (base) and $\text{NH}_4^+$ (its conjugate acid) are the other conjugate pair.

Step 3. Sanity-check the pairing against the reverse reaction: $\text{NH}_4^+$ should be able to donate a proton back to $\text{OH}^-$ , regenerating $\text{NH}_3 + \text{H}_2\text{O}$ — it can, confirming both pairs are correctly matched and that water is behaving as an acid in this particular reaction (it would act as a base instead if paired with a stronger acid like $\text{HCl}$ ).

Step 4. Now re-examine the same reaction through the Lewis lens. In the forward step, the lone pair on nitrogen in $\text{NH}_3$ attacks a proton being released by water, forming a new $N\text{–}H$ bond. Nitrogen's lone pair is the electron-pair donor ( $\text{NH}_3$ = Lewis base), and the proton is the electron-pair acceptor (the proton, momentarily free of water, is the Lewis acid). The classification agrees exactly with Step 1: the Brønsted base is also the Lewis base.

Step 5. State the final answer. Brønsted acid: $\text{H}_2\text{O}$ ; Brønsted base: $\text{NH}_3$ ; conjugate pairs: $\text{H}_2\text{O}/\text{OH}^-$ and $\text{NH}_4^+/\text{NH}_3$ ; and the same electron-pair roles hold under the Lewis definition, with $\text{NH}_3$ as the Lewis base and the transferring proton as the Lewis acid in both classifications.

Competitive stretch

Now classify $\text{BF}_3 + :\text{NH}_3 \rightarrow \text{F}_3\text{B}\leftarrow\text{NH}_3$ . Boron in $\text{BF}_3$ has only three bonding pairs and an empty $p$ orbital, so it can only accept an electron pair — it is the Lewis acid. Nitrogen's lone pair is donated into that empty orbital, so $\text{NH}_3$ is again the Lewis base, forming the coordinate bond $\text{B}\leftarrow\text{N}$ . But no proton moves and no water is involved, so this reaction has no Brønsted acid, no Brønsted base, and no Arrhenius classification at all — only the Lewis definition, being the widest circle, can describe it.

Common traps

Assuming all three theories always agree. They agree for ordinary aqueous acids and bases, but only the Lewis definition can classify reactions with no water and no proton transfer, such as $\text{BF}_3 + \text{NH}_3$ .
Mispairing conjugates. A conjugate acid–base pair must differ by exactly one $H^+$ , with the acid being the species carrying the extra proton — never pair species from opposite sides of the equation that differ by more than one proton.
Forgetting water is amphiprotic. Water is not permanently "the base" or "the acid"; its role depends on what it reacts with, so check the direction of proton transfer in each specific reaction rather than assuming.
Confusing Lewis acidity with ionisation strength. Being a good Lewis acid (a good electron-pair acceptor) is an electronic/structural property in a given reaction, not the same statement as being a strong Arrhenius acid that ionises fully in water.

Ask the tutor: Ask the AI tutor for a rapid-fire set of reactions to classify as Arrhenius, Bronsted-Lowry, and/or Lewis, including a few tricky no-water cases.

If Arrhenius acids are always Bronsted-Lowry acids too, why do textbooks still teach the Arrhenius definition at all

Because it is the simplest starting point and matches almost everything a beginner meets first — aqueous acids and bases like HCl and NaOH. It builds the right intuition about H+ and OH- before the more abstract proton-transfer and electron-pair pictures are introduced. You still need it as the base case; you just need to remember it is a special case, not the full story.

Is a proton, H+, really a Lewis acid on its own, and how does that work if it has no electrons at all

Yes — a bare proton has an empty 1s orbital and nothing to fill it, so by definition it can accept an electron pair, which is exactly what makes it a Lewis acid. When a base's lone pair forms a new bond to that empty orbital, you get the same event chemists call proton transfer in the Bronsted-Lowry picture. The two descriptions are two views of one mechanism, which is precisely why every Bronsted-Lowry reaction is automatically a Lewis reaction too.

How do I quickly tell if a species can act as a Lewis acid without drawing orbitals every time

Look for either an incomplete octet (like boron in BF3, which has only six electrons around it) or a positively charged metal cation (like Al3+), since both signal an empty or accessible low-energy orbital that can accept a lone pair. Conversely, look for an available lone pair (like the nitrogen in NH3 or the oxygen in H2O) to spot a likely Lewis base. It takes practice, but electron-deficient central atoms and cations are the fastest tell for Lewis acids.