Brønsted–Lowry Acids and Bases

In 1923, two chemists working independently, Johannes Brønsted in Copenhagen and Thomas Lowry in Cambridge, published the same idea within months of each other. Instead of defining acids and bases by which ions they produce in water, they defined them by what happens to a single particle: the proton.

That shift sounds small, but it made the theory work in any solvent, or none at all, and explained why substances like ammonia are bases. It’s the definition used throughout most of secondary and university chemistry.

The definitions

• A Brønsted–Lowry acid is a proton donor.
• A Brønsted–Lowry base is a proton acceptor.

“Proton” here means a hydrogen ion, H⁺: a hydrogen atom that has lost its only electron.

An acid–base reaction is therefore a proton transfer. One particle gives up an H⁺, another takes it. You can’t have one without the other.

A first example

Hydrogen chloride dissolving in water:

HCl + H₂O → H₃O⁺ + Cl⁻

• HCl gives a proton → acid
• H₂O accepts the proton → base

That second point surprises people. Water, usually thought of as neutral, is acting as a base here. In Brønsted–Lowry theory, a substance’s role depends on what it’s reacting with.

Ammonia: the case Arrhenius couldn’t handle

NH₃ + H₂O ⇌ NH₄⁺ + OH⁻

• NH₃ accepts a proton → base
• H₂O donates a proton → acid

Ammonia has no hydroxide in it, yet it makes water alkaline. Brønsted–Lowry theory explains this cleanly: ammonia pulls a proton off water, leaving hydroxide behind. The nitrogen’s lone pair of electrons is what grabs the proton.

And here water is the acid. Water can act as either an acid or a base depending on its partner, which makes it amphiprotic (a type of amphoteric behaviour).

Reactions without water

The biggest advantage of the theory is that it doesn’t need water:

HCl(g) + NH₃(g) → NH₄Cl(s)

HCl donates a proton to NH₃. The two colourless gases produce a white smoke of ammonium chloride where they meet. This is a classic demonstration and a genuine acid–base reaction with no solvent involved.

Conjugate acid–base pairs

When an acid loses a proton, what’s left behind can accept a proton back. That leftover particle is the acid’s conjugate base. Likewise, when a base gains a proton, it becomes its conjugate acid.

A conjugate pair always differs by exactly one H⁺.

Look again at the ammonia reaction:

NH₃ + H₂O ⇌ NH₄⁺ + OH⁻

Pair	Acid	Base
Pair 1	NH₄⁺	NH₃
Pair 2	H₂O	OH⁻

Every Brønsted–Lowry reaction has two conjugate pairs: acid₁/base₁ and acid₂/base₂. The forward reaction is acid₂ (H₂O) donating to base₁ (NH₃). The reverse reaction is acid₁ (NH₄⁺) donating to base₂ (OH⁻).

How to find the pairs

1. Identify which reactant loses an H⁺ (the acid) and which gains one (the base).
2. Match each reactant with the product that differs from it by one H⁺.
3. The species with the extra H⁺ in each pair is the acid.

Worked examples

Example 1: HNO₃ + H₂O → H₃O⁺ + NO₃⁻

• Acid: HNO₃; its conjugate base: NO₃⁻
• Base: H₂O; its conjugate acid: H₃O⁺

Example 2: CH₃COOH + H₂O ⇌ CH₃COO⁻ + H₃O⁺

• Acid: CH₃COOH; conjugate base: CH₃COO⁻ (ethanoate)
• Base: H₂O; conjugate acid: H₃O⁺

Example 3: CO₃²⁻ + H₂O ⇌ HCO₃⁻ + OH⁻

• Base: CO₃²⁻; conjugate acid: HCO₃⁻
• Acid: H₂O; conjugate base: OH⁻

This is why sodium carbonate solution is alkaline: the carbonate ion takes protons from water and generates hydroxide.

Example 4: HSO₄⁻ + NH₃ → SO₄²⁻ + NH₄⁺

• Acid: HSO₄⁻; conjugate base: SO₄²⁻
• Base: NH₃; conjugate acid: NH₄⁺

Example 5: H₂O + H₂O ⇌ H₃O⁺ + OH⁻ Water reacting with itself (autoionisation). One molecule acts as an acid, the other as a base. This tiny equilibrium is why pure water has a pH of 7 at 25 °C.

Strength and conjugates: the see-saw

There’s an inverse relationship between the strength of an acid and the strength of its conjugate base.

• Strong acid → very weak conjugate base. HCl gives up its proton completely, so Cl⁻ has almost no tendency to take one back. Chloride is essentially neutral in water.
• Weak acid → noticeably basic conjugate base. Ethanoic acid holds onto its proton fairly tightly, so the ethanoate ion readily takes protons back. That makes sodium ethanoate solution slightly alkaline.

Acid	Strength	Conjugate base	Basic strength
HCl	Very strong	Cl⁻	Negligible
HNO₃	Very strong	NO₃⁻	Negligible
HF	Weak	F⁻	Weak base
CH₃COOH	Weak	CH₃COO⁻	Weak base
NH₄⁺	Very weak	NH₃	Weak base
H₂O	Extremely weak	OH⁻	Strong base

Equilibrium always favours the side with the weaker acid and base. That lets you predict direction: a proton transfer goes mainly towards forming the weaker acid.

Using the theory: why some salts aren’t neutral

Brønsted–Lowry thinking lets you predict the pH of a salt solution without memorising a list. Split the salt into its ions and ask whether either ion is a meaningful acid or base.

• Sodium chloride: Na⁺ doesn’t donate or accept protons; Cl⁻ is the conjugate base of a strong acid, so it’s negligibly basic. Solution: neutral.
• Sodium ethanoate: Na⁺ is a spectator; CH₃COO⁻ is the conjugate base of a weak acid and takes protons from water. Solution: slightly alkaline.
• Ammonium chloride: Cl⁻ is a spectator; NH₄⁺ is the conjugate acid of a weak base and donates protons to water. Solution: slightly acidic.

The Arrhenius definition gives no hint why these salts differ. Brønsted–Lowry makes the answer almost automatic.

Amphiprotic species

Some particles can both donate and accept a proton. Besides water, common examples are:

• HCO₃⁻ (hydrogencarbonate): can become H₂CO₃ or CO₃²⁻
• HSO₄⁻ (hydrogensulfate): can become H₂SO₄ or SO₄²⁻, though in practice it mostly acts as an acid
• H₂PO₄⁻ and HPO₄²⁻ (phosphates): both are important buffers in your cells
• Amino acids: have both an acidic –COOH group and a basic –NH₂ group

The hydrogencarbonate ion is especially important. It’s the main buffer in your blood, mopping up extra acid or extra base to keep blood pH close to 7.4. Read more in buffers explained.

Common mistakes

• Mixing up H and H⁺. A conjugate pair differs by one proton (H⁺), so the charge changes by one as well. NH₃ → NH₄⁺, not NH₄.
• Pairing the wrong species. The conjugate of a reactant is always on the other side of the equation.
• Thinking water is always neutral. In Brønsted–Lowry reactions water is often the acid or the base.
• Saying “strong conjugate base” for strong acids. It’s the opposite: strong acids have extremely weak conjugate bases.

Practice questions

Identify the acid, base, conjugate acid and conjugate base:

1. HF + H₂O ⇌ F⁻ + H₃O⁺
2. HCO₃⁻ + OH⁻ → CO₃²⁻ + H₂O
3. CH₃NH₂ + H₂O ⇌ CH₃NH₃⁺ + OH⁻
4. H₂PO₄⁻ + H₂O ⇌ HPO₄²⁻ + H₃O⁺

Answers:

1. Acid HF, base H₂O, conjugate base F⁻, conjugate acid H₃O⁺.
2. Acid HCO₃⁻, base OH⁻, conjugate base CO₃²⁻, conjugate acid H₂O.
3. Base CH₃NH₂, acid H₂O, conjugate acid CH₃NH₃⁺, conjugate base OH⁻.
4. Acid H₂PO₄⁻, base H₂O, conjugate base HPO₄²⁻, conjugate acid H₃O⁺.

Key takeaways

• Acids donate protons; bases accept them. Every acid–base reaction is a proton transfer.
• The theory works in any solvent and explains bases without hydroxide, such as ammonia.
• Conjugate pairs differ by exactly one H⁺. Each reaction has two pairs.
• The stronger the acid, the weaker its conjugate base.
• Water and ions like HCO₃⁻ are amphiprotic: they can act as either acid or base.
• To go further, compare Brønsted–Lowry with the Arrhenius and Lewis views in acid–base definitions.