Bronsted-Lowry Base
Definition and meaning of Bronsted-Lowry Base in chemistry.
A Bronsted-Lowry base is any species that accepts a proton (H+) from another substance during a chemical reaction. This definition, proposed independently by Johannes Bronsted and Thomas Lowry in 1923, broadened the concept of a base beyond the Arrhenius model, which required a hydroxide ion.
In more detail
Because the definition focuses on proton transfer rather than the presence of a specific ion, it applies to substances dissolved in water and to reactions in nonaqueous or gas-phase settings alike. When a Bronsted-Lowry base accepts a proton, it becomes its conjugate acid, forming a conjugate acid-base pair. The base must have an available lone pair of electrons or a bonding pair capable of forming a new bond to the incoming proton.
Key facts
| Proposed by | J.N. Bronsted and T.M. Lowry, 1923 |
|---|---|
| Defining action | Accepts a proton (H+) |
| Produces | Conjugate acid upon accepting H+ |
| Field | General Chemistry |
In the reaction NH3 + H2O -> NH4+ + OH-, ammonia (NH3) acts as the Bronsted-Lowry base by accepting a proton from water, forming the conjugate acid NH4+.
Frequently asked questions
How does a Bronsted-Lowry base differ from an Arrhenius base?
An Arrhenius base must release OH- ions in water, while a Bronsted-Lowry base simply accepts a proton, so it applies to a wider range of substances and solvents, including ammonia and other species lacking hydroxide groups.
Can water act as a Bronsted-Lowry base?
Yes, water is amphoteric: it acts as a base when it accepts a proton to form H3O+, and as an acid when it donates a proton to form OH-.