Brønsted–Lowry acids and bases

Key structures for this topic — drawn live.

The core definition

The Brønsted–Lowry theory defines acids and bases by what they do with a proton — a hydrogen nucleus, H⁺, which is just a hydrogen atom stripped of its electron. An acid is a proton donor and a base is a proton acceptor. Every Brønsted acid–base reaction is simply the transfer of one H⁺ from the acid to the base.

Notice that being a base requires a place to put the incoming proton, which means a base must have a lone pair of electrons (or occasionally a π bond) available to form the new bond to hydrogen. That lone pair is doing the work.

Conjugate acid–base pairs

Every acid has a conjugate base: what's left after it donates its proton. Every base has a conjugate acid: what it becomes after it accepts a proton. The rule is simple — members of a conjugate pair differ by exactly one H⁺ (one proton and its positive charge).

• HCl (acid) loses H⁺ → Cl⁻ (its conjugate base).
• H₂O (base) gains H⁺ → H₃O⁺ (its conjugate acid).
• H₂O (acid) loses H⁺ → OH⁻ (its conjugate base).

Water shows up as both an acid and a base depending on its partner — a substance that can act either way is called amphoteric.

The curved arrows

In a mechanism, an acid–base reaction is drawn with two curved arrows. The first arrow starts at the base's lone pair and points to the acidic hydrogen — this forms the new bond to H. The second arrow starts at the H–X bond of the acid and points onto X, giving those bonding electrons to the atom that's left behind.

The result: the base is now bonded to H (it became the conjugate acid), and the acid's former partner atom keeps both electrons and picks up a negative charge (it became the conjugate base). Curved arrows always follow the electrons, never the proton itself — the proton moves because electrons pushed toward it and away from it.

Strong acids and weak conjugate bases

There's an inverse relationship you'll use constantly: the stronger the acid, the weaker its conjugate base, and vice versa. A strong acid like HCl gives up its proton eagerly precisely because the resulting Cl⁻ is very stable and has little tendency to grab the proton back — so Cl⁻ is a weak base.

Conversely, a weak acid such as water holds its proton tightly, so its conjugate base, hydroxide (OH⁻), is a relatively strong base. This is why acid–base reactions always run in the direction that forms the weaker acid and the weaker base — the more stable side is favored at equilibrium.

Identifying acid, base, and conjugates in a reaction

To label the four species in any Brønsted reaction, follow the hydrogen:

1. Find the species that loses an H going left to right — that reactant is the acid.
2. The species that gains that H is the base.
3. On the product side, the acid minus its H is the conjugate base.
4. The base plus its new H is the conjugate acid.

Check your work: each labeled pair should differ by one proton and one unit of charge.

A worked example: a carboxylic acid

Consider acetic acid reacting with hydroxide: CH₃COOH + OH⁻ → CH₃COO⁻ + H₂O. Acetic acid donates the proton from its O–H group, so it is the acid; hydroxide accepts it, so OH⁻ is the base. The products are acetate, CH₃COO⁻ (the conjugate base of acetic acid), and water, H₂O (the conjugate acid of hydroxide).

Carboxylic acids are good Brønsted acids largely because their conjugate base, the carboxylate, is stabilized by resonance — the negative charge is spread over two equivalent oxygens. That stability is what makes the acid willing to donate its proton in the first place, which ties directly into how we compare acid strength using pK_a.