The Brønsted–Lowry definition, formulated in 1923, independently by Johannes Nicolaus Brønsted in Denmark and Martin Lowry in England, is based upon the idea of protonation of bases through the deprotonation of acids – that is, the ability of acids to "donate" hydrogen ions () otherwise known as protons to bases, which "accept" them.
An acid–base reaction is, thus, the removal of a hydrogen ion from the acid and its addition to the base. The removal of a hydrogen ion from an acid produces its conjugate base, which is the acid with a hydrogen ion removed. The reception of a proton by a base produces its conjugate acid, which is the base with a hydrogen ion added.
Unlike the previous definitions, the Brønsted–Lowry definition does not refer to the formation of salt and solvent, but instead to the formation of conjugate acids and conjugate bases, produced by the transfer of a proton from the acid to the base. In this approach, acids and bases are fundamentally different in behavior from salts, which are seen as electrolytes, subject to the theories of Debye, Onsager, and others. An acid and a base react not to produce a salt and a solvent, but to form a new acid and a new base. The concept of neutralization is thus absent. Brønsted–Lowry acid–base behavior is formally independent of any solvent, making it more all-encompassing than the Arrhenius model. The calculation of pH under the Arrhenius model depended on alkalis (bases) dissolving in water (aqueous solution). The Brønsted–Lowry model expanded what could be pH tested using insoluble and soluble solutions (gas, liquid, solid).
The general formula for acid–base reactions according to the Brønsted–Lowry definition is:
where HA represents the acid, B represents the base,  represents the conjugate acid of B, and  represents the conjugate base of HA.
For example, a Brønsted–Lowry model for the dissociation of hydrochloric acid (HCl) in aqueous solution would be the following:
The removal of  from the  produces the chloride ion, , the conjugate base of the acid. The addition of  to the  (acting as a base) forms the hydronium ion, , the conjugate acid of the base.
Water is amphoteric that is, it can act as both an acid and a base. The Brønsted–Lowry model explains this, showing the dissociation of water into low concentrations of hydronium and hydroxide ions:
This equation is demonstrated in the image below:
Here, one molecule of water acts as an acid, donating an  and forming the conjugate base, , and a second molecule of water acts as a base, accepting the  ion and forming the conjugate acid, .
As an example of water acting as an acid, consider an aqueous solution of pyridine, .
In this example, a water molecule is split into a hydrogen ion, which is donated to a pyridine molecule, and a hydroxide ion.
In the Brønsted–Lowry model, the solvent does not necessarily have to be water, as is required by the Arrhenius Acid–Base model. For example, consider what happens when acetic acid, , dissolves in liquid ammonia.
An  ion is removed from acetic acid, forming its conjugate base, the acetate ion, . The addition of an  ion to an ammonia molecule of the solvent creates its conjugate acid, the ammonium ion, .
The Brønsted–Lowry model calls hydrogen-containing substances (like ) acids. Thus, some substances, which many chemists considered to be acids, such as  or , are excluded from this classification due to lack of hydrogen. Gilbert N. Lewis wrote in 1938, "To restrict the group of acids to those substances that contain hydrogen interferes as seriously with the systematic understanding of chemistry as would the restriction of the term oxidizing agent to substances containing oxygen." Furthermore,  and  are not considered Brønsted bases, but rather salts containing the bases  and .