Lewis acid–base theory
/ LOO-iss /
Forget protons for a moment and think about who is sharing what. In the Lewis picture, an acid–base reaction is one partner offering a pair of electrons and another partner reaching out to accept them — like one hand holding out a key and another taking it to form a connection. This is a broader, electron-centred way to see acids and bases.
A Lewis base is an electron-pair donor; a Lewis acid is an electron-pair acceptor. They join by forming a new bond in which the base supplies both shared electrons. This definition contains the proton-based Brønsted–Lowry view as a special case (a proton is just one kind of electron-pair acceptor), but it reaches much further: it includes reactions with no proton at all, such as a metal ion bonding to ammonia.
Lewis theory matters because it unifies a vast range of chemistry — catalysis, complex-ion formation, and much of organic reaction mechanism — under one simple idea of electron-pair sharing. The trade-off is that it is so general it no longer maps neatly onto the familiar pH scale; for ordinary water solutions, the Brønsted–Lowry proton view is usually the more practical tool.
Ammonia (NH3) bonds to boron trifluoride (BF3) even though no proton moves: nitrogen donates its lone electron pair into boron's empty slot. NH3 is the Lewis base, BF3 the Lewis acid.
Lewis acid–base = electron-pair acceptor meets electron-pair donor; no proton required.
The three definitions nest: every Brønsted–Lowry acid–base reaction is also a Lewis one, but not the reverse. Choose the narrowest theory that still covers your reaction — usually Brønsted–Lowry for water chemistry, Lewis when electrons, not protons, move.