Brønsted-Lowry Theory
The Brønsted-Lowry theory, proposed independently by Johannes Nicolaus Brønsted and Thomas Martin Lowry in 1923, defines an acid as a proton (H⁺) donor and a base as a proton acceptor, expanding the earlier Arrhenius definition beyond aqueous solutions. This broader definition explains acid-base behaviour in non-aqueous solvents and introduces the concept of conjugate acid-base pairs, where the product of an acid after donating a proton is its conjugate base, and vice versa. The Brønsted-Lowry model is the most widely used framework for acid-base chemistry in biology, organic chemistry, and analytical chemistry.
Comparison of Acid-Base Theories
| Theory | Acid Definition | Base Definition | Limitation |
|---|---|---|---|
| Arrhenius (1884) | Produces H⁺ in water | Produces OH⁻ in water | Only applies in aqueous solutions |
| Brønsted-Lowry (1923) | Proton (H⁺) donor | Proton (H⁺) acceptor | Requires transferable proton |
| Lewis (1923) | Electron pair acceptor | Electron pair donor | Broadest; less intuitive for pH work |
| Lux-Flood (1939) | Oxide ion (O²⁻) acceptor | Oxide ion donor | Applies mainly to high-T melts |
| Solvent system theory | Gives solvent cation | Gives solvent anion | Limited to specific solvents |
Interactive Tools
Wikimedia Commons, CC BY-SA
Related Terms
Acid (Chemistry)
An acid is a substance that donates protons (hydrogen ions, H⁺) to another substance or accepts electron pairs, resulting in a sour taste, ability to turn blue litmus red, and a pH below 7 in aqueous solution. Acids play a fundamental role in chemical reactions, biological processes, and industrial applications ranging from digestion to manufacturing. Common examples include hydrochloric acid (HCl), sulfuric acid (H₂SO₄), and acetic acid (CH₃COOH) found in vinegar.
Base (Chemistry)
A base is a substance that accepts protons (H⁺ ions) from an acid or donates electron pairs, producing hydroxide ions (OH⁻) in aqueous solution, a bitter taste, slippery feel, and a pH above 7. Bases neutralise acids in reactions that form salt and water, making them essential in biological systems such as blood buffering, as well as in industrial processes like soap making. Common examples include sodium hydroxide (NaOH), ammonia (NH₃), and sodium bicarbonate (NaHCO₃).
Acid Dissociation Constant (Ka)
The acid dissociation constant (Ka) is the equilibrium constant for the ionisation of an acid in water, quantifying the extent to which an acid donates its proton to water at a given temperature. A large Ka (> 1) indicates a strong acid that dissociates almost completely, while a small Ka (< 10⁻³) indicates a weak acid with limited dissociation. Ka values are fundamental in predicting reaction directions, calculating pH of weak acid solutions, and designing buffer systems.
Named after Johannes Nicolaus Brønsted (Danish, 1879–1947) and Thomas Martin Lowry (British, 1874–1936), who independently published the proton-transfer theory in 1923. "Brønsted" is a Danish surname; "Lowry" is of English origin. The theory refined "acid" from Latin "acidus" (sour) and "base" from Latin "basis" (foundation) in the context of proton transfer.