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Three progressively broader definitions classify acids and bases. The Arrhenius definition (narrowest): acids produce H⁺ in water, bases produce OH⁻. The Brønsted-Lowry definition: acids are proton (H⁺) donors, bases are proton acceptors — this works in any solvent and introduces conjugate acid-base pairs (an acid donates a proton to become its conjugate base, and vice versa). The Lewis definition (broadest): acids are electron-pair acceptors, bases are electron-pair donors — encompassing reactions with no proton transfer at all, such as BF₃ accepting a lone pair from NH₃. Each broader definition includes all reactions classified by the narrower one.
Classify the same reaction under all three definitions where possible, then find examples that work under Lewis but not Brønsted-Lowry (e.g., metal-ligand coordination). Practice identifying conjugate pairs: every Brønsted-Lowry reaction has exactly two conjugate pairs.
The simplest way to think about acids and bases starts with water. The Arrhenius definition says an acid is any substance that produces H⁺ ions when dissolved in water, and a base produces OH⁻ ions. HCl dissolves and releases H⁺; NaOH dissolves and releases OH⁻. This works well for straightforward aqueous reactions, but it immediately runs into limits. What about ammonia, NH₃, which makes solutions basic without containing any OH⁻ in its formula? And what about reactions that happen in solvents other than water, or with no solvent at all? You need a broader framework.
The Brønsted-Lowry definition solves this by focusing on proton transfer rather than what dissolves in water. An acid is a proton (H⁺) donor; a base is a proton acceptor. When HCl reacts with NH₃, HCl donates a proton to NH₃ — HCl is the acid, NH₃ is the base. This definition introduces a powerful concept: conjugate pairs. After HCl donates its proton, it becomes Cl⁻, which is HCl's conjugate base. After NH₃ accepts the proton, it becomes NH₄⁺, which is NH₃'s conjugate acid. Every Brønsted-Lowry reaction produces exactly two conjugate pairs. From your understanding of covalent bonding, you can see why this works — the proton transfer involves breaking one covalent bond (H–Cl) and forming another (N–H). The strength of these bonds determines how readily the transfer occurs.
The Lewis definition takes one more step outward. Instead of tracking protons, it tracks electron pairs. A Lewis acid accepts an electron pair; a Lewis base donates one. This is the broadest definition because it captures reactions with no proton involved at all. When BF₃ reacts with NH₃, boron has an empty orbital that accepts the lone pair on nitrogen — BF₃ is the Lewis acid, NH₃ is the Lewis base, and a new coordinate covalent bond forms. Metal ions in solution act as Lewis acids when water molecules donate lone pairs to them during hydration. None of these involve proton transfer, yet they follow the same underlying logic of electron-pair sharing that you learned in covalent bonding.
The three definitions are nested like concentric circles: every Arrhenius acid-base reaction is also a Brønsted-Lowry reaction, and every Brønsted-Lowry reaction is also a Lewis reaction — but not the reverse. In practice, chemists default to Brønsted-Lowry for most aqueous chemistry and reach for the Lewis definition when dealing with coordination chemistry, organic reaction mechanisms, or any scenario where protons are not the central players. The key insight is that these are not competing theories but progressively wider lenses for the same fundamental phenomenon: the movement of electron density between species.