Brønsted-Lowry (Bronsted-Lowry) acids and bases

A Brønsted-Lowry acid is a proton donor: a substance that gives a hydrogen ion, H⁺, to another substance. A Brønsted-Lowry base is a proton acceptor. Every acid-base reaction in this theory is a proton transfer from the acid to the base. A hydrogen atom that has lost its only electron is just a proton, which is why H⁺ and "proton" mean the same thing here.

What is a Brønsted-Lowry acid?

AQA A-level Chemistry states it in three lines: an acid is a proton donor, a base is a proton acceptor, and acid-base equilibria involve the transfer of protons. OCR uses the same definitions and adds conjugate acid-base pairs.

Take hydrogen chloride dissolving in water:

HCl + H₂O → H₃O⁺ + Cl⁻

• HCl gives a proton to water, so HCl is the acid.
• Water takes the proton, so here water is the base.
• The product H₃O⁺ is the oxonium ion (also called the hydronium ion). It is what "H⁺(aq)" really means in an aqueous solution.

Brønsted-Lowry bases

A base needs a lone pair of electrons to accept the proton. Ammonia uses the lone pair on its nitrogen:

NH₃ + H₂O ⇌ NH₄⁺ + OH⁻

This time water gives a proton, so water is the acid and ammonia is the base. The OH⁻ ions formed make the solution alkaline.

Conjugate acid-base pairs

When an acid gives away a proton, what is left can take a proton back, so it is a base: the conjugate base of that acid. When a base accepts a proton, it becomes its conjugate acid. The two members of a conjugate pair differ by exactly one H⁺.

To find the pairs in any equation, look for two species, one on each side, whose formulas differ by one H and one positive charge.

Water and other amphoteric species

In the tables above, water is a base with HCl but an acid with ammonia. A substance that can either donate or accept a proton is called amphiprotic (amphoteric is also used). Water even transfers protons between its own molecules, to a very small extent:

H₂O + H₂O ⇌ H₃O⁺ + OH⁻

OpenStax notes that at 25 °C only about two in every billion water molecules are ionised this way. This equilibrium leads to the ionic product of water, K_w. The hydrogencarbonate ion, HCO₃⁻, is another amphiprotic species: it can lose a proton to become CO₃²⁻ or gain one to become H₂CO₃.

Brønsted-Lowry vs Arrhenius

Johannes Brønsted and Thomas Lowry proposed their more general description in 1923. It explains reactions that happen without water, such as hydrogen chloride gas meeting ammonia gas to form solid ammonium chloride: HCl(g) + NH₃(g) → NH₄Cl(s). HCl donates a proton to NH₃, even though no water is present and no OH⁻ ions form.

Strong and weak acids

• Strong acids, such as HCl and HNO₃, donate their protons completely in water: they fully dissociate.
• Weak acids, such as ethanoic acid, CH₃COOH, only partly dissociate. The reaction is an equilibrium, so most of the acid molecules stay intact. Their strength is measured by K_a; see the pKa guide.

OCR also asks about monobasic, dibasic and tribasic acids: how many protons each molecule can donate. HCl is monobasic, sulfuric acid, H₂SO₄, is dibasic and phosphoric acid, H₃PO₄, is tribasic.

Worked examples

Example 1: identify the acid, base and pairs

For HNO₂ + H₂O ⇌ NO₂⁻ + H₃O⁺, name the Brønsted-Lowry acid and base and the two conjugate pairs.

• HNO₂ loses a proton to become NO₂⁻: HNO₂ is the acid, and NO₂⁻ is its conjugate base.
• H₂O gains a proton to become H₃O⁺: H₂O is the base, and H₃O⁺ is its conjugate acid.
• The pairs are HNO₂/NO₂⁻ and H₃O⁺/H₂O.

Example 2: water as the acid

Methylamine reacts with water: CH₃NH₂ + H₂O ⇌ CH₃NH₃⁺ + OH⁻. Which species is the Brønsted-Lowry acid?

• Water gives a proton to methylamine and becomes OH⁻, so water is the acid. Methylamine is the base, using the lone pair on its nitrogen, and CH₃NH₃⁺ is its conjugate acid.

Example 3: pH of a strong acid

Find the pH of 0.050 mol dm⁻³ nitric acid.

1. HNO₃ is a strong, monobasic acid, so it donates all its protons: [H⁺] = 0.050 mol dm⁻³.
2. pH = −log₁₀[H⁺] = −log₁₀(0.050) = 1.30.
3. The pH Calculator handles strong and weak acids, and the how to calculate pH guide has more examples.

Common mistakes

• Saying an acid donates hydrogen. It donates a proton, H⁺, not a hydrogen atom or H₂.
• Pairing the acid with the wrong species. A conjugate pair differs by exactly one H⁺: HCl/Cl⁻, not HCl/H₃O⁺.
• Thinking water is always neutral in these reactions. It can be the acid or the base, depending on what it reacts with.
• Confusing strong with concentrated. Strong means fully dissociated; concentrated means a lot of acid per dm³.
• Using H⁺ alone in water. In aqueous solution the proton is carried as H₃O⁺, though H⁺(aq) is accepted shorthand.

Practise with a short timed paper

Brønsted-Lowry acids and bases are in the second year of A-level Chemistry: AQA 3.1.12 (A-level only) and OCR 5.1.3. The daily A-level Chemistry: Year 2 paper is a short timed paper of 8 questions, with a new one every day. For acids at GCSE, see neutralisation.

Read next

Chemistry · Acids and basespH calculator: strong and weak acids, bases and buffersFree pH calculator: the pH of strong and weak acids and bases from molarity, buffers by Henderson-Hasselbalch, and Ka from a measured pH, with steps.