Acid–Base Definitions

Three ways to define an acid — and why chemists reach for a different one depending on the reaction in front of them.

High schoolIntro Gen ChemUni Year 1
⏱️ About 16 min

Vinegar tastes sour, soap feels slippery, and a squeeze of lemon fizzes on baking soda. For centuries 'acid' and 'base' were just lists of behaviours. The breakthrough was realising all of it comes down to one tiny thing changing hands: a proton.

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The big idea: There isn't one definition of an acid — there are three, each broader than the last. Arrhenius talks about H⁺ and OH⁻ in water; Brønsted–Lowry is about donating and accepting protons; Lewis is about sharing electron pairs. Pick the one that fits your reaction.
🎯 By the end, you'll be able to
  • State the Arrhenius, Brønsted–Lowry and Lewis definitions of acids and bases
  • Identify the proton donor and acceptor in a Brønsted–Lowry reaction
  • Write the conjugate base of an acid (and conjugate acid of a base)
  • Explain why a substance like NH₃ counts as a base even without OH⁻
📎 Helpful to know first

Arrhenius: the water-only picture

The first modern definition, from Svante Arrhenius, is the one most people meet first. In water:

  • An acid releases hydrogen ions, H⁺ (really H₃O⁺).
  • A base releases hydroxide ions, OH⁻.

Hydrochloric acid fits perfectly — drop it in water and it hands over H⁺. Sodium hydroxide fits too — it releases OH⁻. Neutralisation is then just H⁺ meeting OH⁻ to make water. Simple and useful… as long as you stay in water and stick to OH⁻ bases.

\[ \ce{HCl -> H+ + Cl-} \qquad \ce{NaOH -> Na+ + OH-} \]
The Arrhenius view: acids give H⁺, bases give OH⁻ in aqueous solution.

Brønsted–Lowry: it's about protons, not water

Ammonia (NH₃) is clearly basic — it turns litmus blue — yet it contains no OH⁻ to give away. Arrhenius can't explain that. Brønsted and Lowry fixed it by shifting the focus from what's released to what's transferred:

  • An acid is a proton (H⁺) donor.
  • A base is a proton acceptor.

Now ammonia makes sense: it accepts a proton from water, leaving OH⁻ behind. No hydroxide needed in the starting material — the base simply grabs a proton.

\[ \ce{NH3 + H2O <=> NH4+ + OH-} \]
Ammonia accepts a proton from water (water is the acid here), producing OH⁻.
🔑 Conjugate acid–base pairs
When an acid donates its proton, what's left is its conjugate base. When a base accepts a proton, it becomes its conjugate acid. They differ by exactly one H⁺. In HCl + H₂O → H₃O⁺ + Cl⁻: HCl/Cl⁻ are a pair, and H₂O/H₃O⁺ are the other pair.
✨ Water plays both roles
Notice water was an acid toward ammonia but a base toward HCl. A substance that can donate or accept a proton is amphoteric. Which role it plays depends on its partner — there's nothing special baked into the molecule that fixes it as one or the other.
📝 Worked example: In the reaction HF + H₂O ⇌ H₃O⁺ + F⁻, identify each Brønsted–Lowry acid, base, and their conjugates.
  1. Look for the proton (H⁺) transfer. HF loses an H⁺, so HF is the acid (donor).
  2. H₂O gains that H⁺ to become H₃O⁺, so H₂O is the base (acceptor).
  3. Remove the donated proton from HF → F⁻, so F⁻ is the conjugate base of HF.
  4. Add the accepted proton to H₂O → H₃O⁺, so H₃O⁺ is the conjugate acid of water.
✓ Acid HF / conjugate base F⁻; base H₂O / conjugate acid H₃O⁺.

Lewis: the widest lens of all

Some reactions look acid–base in every way but involve no proton at all. Gilbert Lewis went one level deeper — to the electrons that a proton would have bonded to:

  • A Lewis acid accepts a pair of electrons.
  • A Lewis base donates a pair of electrons.

A proton (H⁺) is just one example of a Lewis acid — an empty orbital looking for electrons. But so is BF₃, which has no proton to give. Every Brønsted acid is a Lewis acid, but not every Lewis acid is a Brønsted acid. The definitions nest inside each other.

✨ Three definitions, nested
Arrhenius ⊂ Brønsted–Lowry ⊂ Lewis. Each is a special case of the next. You don't abandon the earlier ones — you reach for the narrowest definition that still captures your reaction, because narrower usually means simpler.

Check your understanding

1. Ammonia (NH₃) is a base even though it contains no OH⁻. Which definition explains this best?
Brønsted–Lowry defines a base as a proton acceptor. NH₃ accepts a proton from water to form NH₄⁺, leaving OH⁻ — no hydroxide needed in the starting material.
2. What is the conjugate base of the acid H₂SO₄ after it donates one proton?
A conjugate base is the acid minus one H⁺. Remove one proton from H₂SO₄ and you get HSO₄⁻ (it can lose a second proton later to become SO₄²⁻).
3. In HCl + H₂O → H₃O⁺ + Cl⁻, what role does water play?
Here water accepts the proton from HCl to become H₃O⁺, so it acts as the base. Water is amphoteric — with ammonia it was the acid instead.
🔎 Related lessons
Next: pH, pOH & Strong Acids/Bases Strong vs Weak Acids & Bases Dynamic Equilibrium & K
✅ Key takeaways
  • Arrhenius: acids release H⁺, bases release OH⁻ — but only in water.
  • Brønsted–Lowry: acids donate protons, bases accept them (explains NH₃).
  • Conjugate pairs differ by exactly one H⁺: HA/A⁻ and B/BH⁺.
  • Lewis: acids accept an electron pair, bases donate one — the broadest view.
  • The three definitions nest: Arrhenius ⊂ Brønsted–Lowry ⊂ Lewis.
➡️ Now that 'acid' means 'proton donor', the obvious next question is: how much H⁺ is actually floating around? That amount is what the pH scale measures — and it's where we head next.
Want to test yourself on this? Try the Chemistry practice test →