What atoms become.

Lemon juice tastes sour — that's an acid at work. Soap feels slippery — that's a base. They're opposites: mix the right amounts and they cancel each other out into something neutral, like plain water. The pH number tells you where a liquid lands — below 7 is the sour side, above 7 is the soapy side, and 7 is the middle.

In water, an acid releases hydrogen ions (H⁺ — a hydrogen atom's bare nucleus) and a base soaks them up. pH counts those ions in a solution. But chemistry class uses a second number too: pKa, which describes the substance itself — how eager it is to hand hydrogen away. Lower pKa, stronger acid. pH is about the glass of liquid in front of you; pKa is about the chemical's personality. The lists below use pKa, because that's what PubChem publishes.

The Brønsted–Lowry picture: an acid donates a proton (H⁺), a base accepts one — and every acid that gives its proton away leaves behind a conjugate base, so the two roles are permanently paired. pKa is the pH at which half of the substance has let go of its proton. Water itself splits into ions in tiny amounts (2 H₂O ⇌ H₃O⁺ + OH⁻), which is why pure water sits at pH 7 at 25 °C — and since water can either donate or accept a proton depending on its partner, it's amphoteric: the textbook both-sides player. Some acids give protons in stages — see phosphoric acid's card above, whose published constants come as pK1, pK2, pK3.

Three definitions, in widening order. Arrhenius: produces H⁺ (acid) or OH⁻ (base) in water — the narrowest, water-only view. Brønsted–Lowry: proton transfer, in any solvent. Lewis: electron-pair acceptor (acid) or donor (base) — no hydrogen required at all, which is how bare metal ions count as acids. pKa = −log₁₀ Ka, the dissociation equilibrium constant. And a real limit of the ruler: in water, every acid stronger than H₃O⁺ looks equally strong (the leveling effect) — which is why strong acids like hydrochloric often carry no meaningful aqueous pKa, and why "is X an acid?" genuinely has no context-free answer.