Bronsted-Lowry Acids and Bases

Calling an acid 'something sour' describes a side effect, not the chemistry. The Bronsted-Lowry definition instead describes the mechanism: an acid is a species that donates a proton (H+), and a base is a species that accepts one. This works for every proton-transfer reaction, even ones you would never taste, like ammonia reacting with HCl gas. Define acids by what they do, not by how they happen to affect your tongue, and the definition never fails you.

Watch what happens when HCl hits water. The H-Cl bond breaks, and that proton does not go free. A water molecule uses a lone pair on its oxygen to grab the proton the instant it's released, forming H_3O+, hydronium, while the leftover Cl- drifts off as chloride. This is why you'll often see H+ written in equations as shorthand, but the real species in solution is hydronium. A bare proton is just a nucleus with no electron cloud around it, so its charge density is enormous; any nearby lone pair swallows it essentially instantly. That is also how proton transfer hops through water generally: one molecule's lone pair captures a proton from a neighbor, over and over, which is why acid-base reactions in water are fast.

Water itself is amphoteric: it can donate a proton (acting as an acid) or accept one (acting as a base), depending on what it meets. This dual identity is what lets water react with itself in autoionization, and it sets up everything in this unit, from the pH scale to conjugate acid-base pairs to how buffers resist pH change.