Bronsted-Lowry acid-base definition

The Arrhenius idea was too tied to water. In 1923 Johannes Bronsted and Thomas Lowry, working independently, offered a cleaner picture: forget about water and hydroxide, and focus on the proton itself. An acid is anything that can DONATE a proton (an H+); a base is anything that can ACCEPT a proton. Acid-base chemistry becomes, at its heart, a proton changing hands.

This shifts the spotlight nicely. Ammonia is now obviously a base — it has a lone pair that grabs a proton to become NH4+, no hydroxide required: NH3 + H2O gives NH4+ + OH-, where water acted as the acid by handing over a proton. Crucially, every acid that donates a proton turns into a base (its conjugate base), and every base that accepts one becomes an acid (its conjugate acid). A reaction is always a competition between two such pairs, and the proton ends up on the stronger base. Water itself can be either: it gives a proton to ammonia but takes one from HCl. We call such double-natured species amphiprotic.

Bronsted-Lowry is the workhorse definition for anything involving H+ — pKa tables, buffers, the relative strengths of inorganic acids like HF, HNO3, and H2SO4, and proton chemistry in non-aqueous solvents such as liquid ammonia. It is broader than Arrhenius (it works in any solvent and even with no solvent) but still narrower than Lewis, because it is fundamentally about the proton. Reactions with no proton transfer at all, like BF3 grabbing an electron pair from NH3, lie outside it.