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Brønsted–Lowry Acid–Base Equilibria



Acids & Bases

AQA Content

(See Specification)

Specification Notes

An acid is a proton donor.
A base is a proton acceptor.
Acid–base equilibria involve the transfer of protons.


Bronsted acids and bases

The theory defines an acid as a substance that donates a proton (H+) and a base as a substance that accepts a proton.

Examples of Bronsted acids include hydrochloric acid (HCl). The hydrogen chloride gives a hydrogen ion (a proton) to a water molecule. A co-ordinate (dative covalent) bond is formed between a lone pair on the oxygen and the hydrogen forming hydroxonium ions, H3O+. It is actually the hydroxonium ion that is the acid.- it gives a proton to the hydroxide ion to give water.

Other Bronsted Acids include sulfuric acid (H2SO4), and acetic acid (CH3COOH). In aqueous solution, they all release a proton, or hydrogen ion, into the water. For example, sulfuric acid (H2SO4) dissociates to form H+ and HSO4-.

Examples of Bronsted bases include ammonia (NH3), water (H2O), and ethanol (C2H5OH). Bases accept a proton from an acid, forming a new compound. For example, when ammonia (NH3) reacts with hydrochloric acid (HCl), it forms a new compound called ammonium chloride (NH4Cl).

Water (H2O) can also act as a Bronsted base by accepting a proton from an acid, forming hydroxonium (H3O+) ions.


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