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The Arrhenius definition of acid and base is limited to aqueous (that is, water) solutions. Although this is useful because water is a common solvent, it is limited to the relationship between the H⁺ ion and the OH⁻ ion. What would be useful is a more general definition that would be more applicable to other chemical reactions and, importantly, independent of H₂O.

In 1923, Danish chemist Johannes Brønsted and English chemist Thomas Lowry independently proposed new definitions for acids and bases, ones that focus on proton transfer. A Brønsted-Lowry acid is any species that can donate a proton (H⁺) to another molecule. A Brønsted-Lowry base is any species that can accept a proton from another molecule. In short, a Brønsted-Lowry acid is a proton donor (PD), while a Brønsted-Lowry base is a proton acceptor (PA).

It is easy to see that the Brønsted-Lowry definition covers the Arrhenius definition of acids and bases. Consider the prototypical Arrhenius acid-base reaction:

[latex]\begin{array}{ccccc} \ce{H+(aq)}&+&\ce{OH-(aq)}&\rightarrow&\ce{H2O(\ell)} \\ \\ \text{(acid)}&&\text{(base)}&& \end{array}[/latex]

The acid species and base species are marked. The proton, however, is (by definition) a proton donor (labelled PD), while the OH⁻ ion is acting as the proton acceptor (labelled PA):

[latex]\begin{array}{ccccc} \ce{H+(aq)}&+&\ce{OH-(aq)}&\rightarrow&\ce{H2O(\ell)} \\ \\ \text{(PD)}&&\text{(PA)}&& \end{array}[/latex]

The proton donor is a Brønsted-Lowry acid, and the proton acceptor is the Brønsted-Lowry base:

[latex]\begin{array}{ccccc} \ce{H+(aq)}&+&\ce{OH-(aq)}&\rightarrow&\ce{H2O(\ell)} \\ \\ \text{(BL acid)}&&\text{(BL base)}&& \end{array}[/latex]

Thus H⁺ is an acid by both definitions, and OH⁻ is a base by both definitions.

Ammonia (NH₃) is a base even though it does not contain OH⁻ ions in its formula. Instead, it generates OH⁻ ions as the product of a proton-transfer reaction with H₂O molecules; NH₃ acts like a Brønsted-Lowry base, and H₂O acts like a Brønsted-Lowry acid:

[latex]\begin{array}{ccccc} \chemfig{H-\Lewis{6:,N}(-[:90]H)-H}&+&\chemfig{H-[:30]O-[:-30]H}&\rightarrow&\chemfig{H-N^+(-[:90]H)(-[:-90]H)-H}+\chemfig{^{-}O-H} \\ \text{proton acceptor}&&\text{proton donor}&& \end{array}[/latex]

A reaction with water is called hydrolysis; we say that NH₃ hydrolyzes to make NH₄⁺ ions and OH⁻ ions.

Even the dissolving of an Arrhenius acid in water can be considered a Brønsted-Lowry acid-base reaction. Consider the process of dissolving HCl(g) in water to make an aqueous solution of hydrochloric acid. The process can be written as follows:

HCl(g) + H₂O(ℓ) → H₃O⁺(aq) + Cl⁻(aq)

HCl(g) is the proton donor and therefore a Brønsted-Lowry acid, while H₂O is the proton acceptor and a Brønsted-Lowry base. These two examples show that H₂O can act as both a proton donor and a proton acceptor, depending on what other substance is in the chemical reaction. A substance that can act as a proton donor or a proton acceptor is called amphiprotic. Water is probably the most common amphiprotic substance we will encounter, but other substances are also amphiprotic.

Observe the reaction between NH₃ and H₂O:

[latex]\begin{array}{ccccc} \chemfig{H-\Lewis{6:,N}(-[:90]H)-H}&+&\chemfig{H-[:30]O-[:-30]H} &\rightarrow & \chemfig{H-N^+(-[:90]H)(-[:-90]H)-H}+\chemfig{O-H^{+}}} \end{array}[/latex]

The chemical reaction does not go to completion; rather, the reverse process occurs as well, and eventually the two processes cancel out any additional change. At this point, we say the chemical reaction is at equilibrium. Both processes still occur, but any net change by one process is countered by the same net change by the other process; it is a dynamic, rather than a static, equilibrium. Because both reactions are occurring, it makes sense to use a double arrow instead of a single arrow:

[latex]\begin{array}{ccccc} \chemfig{H-\Lewis{6:,N}(-[:90]H)-H}&+&\chemfig{H-[:30]O-[:-30]H} &\rightleftharpoons & \chemfig{H-N^+(-[:90]H)(-[:-90]H)-H}+\chemfig{O-H^{-}}} \end{array}[/latex]

What do you notice about the reverse reaction? The NH₄⁺ ion is donating a proton to the OH⁻ ion, which is accepting it. This means that the NH₄⁺ ion is acting as the proton donor, or Brønsted-Lowry acid, while OH⁻ ion, the proton acceptor, is acting as a Brønsted-Lowry base. The reverse reaction is also a Brønsted-Lowry acid base reaction:

[latex]\begin{array}{ccccc} \chemfig{H-\Lewis{6:,N}(-[:90]H)-H}&+&\chemfig{H-[:30]O-[:-30]H} &\rightleftharpoons & \chemfig{H-N^+(-[:90]H)(-[:-90]H)-H}+\chemfig{O-H^{-}}} \end{array}[/latex]

This means that both reactions are acid-base reactions by the Brønsted-Lowry definition. If you consider the species in this chemical reaction, two sets of similar species exist on both sides. Within each set, the two species differ by a proton in their formulas, and one member of the set is a Brønsted-Lowry acid, while the other member is a Brønsted-Lowry base. These sets are marked here:

The two sets — NH₃/NH₄⁺ and H₂O/OH⁻ — are called conjugate acid-base pairs. We say that NH₄⁺ is the conjugate acid of NH₃, OH⁻ is the conjugate base of H₂O, and so forth. Every Brønsted-Lowry acid-base reaction can be labelled with two conjugate acid-base pairs.