Søren Sørensen, a Danish biochemist working at the Carlsberg Laboratory in Copenhagen, defined pH in 1909 — initially writing it as p_H — to express the acidity of beer wort during fermentation. The notation he chose was deliberately mathematical: pH = −log₁₀[H⁺], the negative base-10 logarithm of the hydrogen-ion concentration. The choice was motivated by practical convenience: hydrogen-ion concentrations span enormous ranges, and logarithms compress them. But it turned out to match the way biological systems care about acidity, where most enzymes respond logarithmically to H⁺. The brewing chemist's bookkeeping had accidentally found the scale on which life itself keeps its books.
Acids donate protons; bases accept them. The Brønsted-Lowry definition (1923) recasts the chemistry of acidity as a chemistry of transfer — every acid-base reaction is a proton changing hands, and the medium that catches it is almost always water. Water itself is mildly amphoteric: it autoionizes into a tiny equilibrium of H₃O⁺ and OH⁻, with the product [H⁺][OH⁻] fixed at 10⁻¹⁴ at room temperature. Pure water sits at pH 7; below that is acidic, above is basic, and the practical scale runs from battery acid near zero to drain cleaner near fourteen. Most chemistry of interest happens with weak acids and bases — those that only partially ionize — and a single equation describes how they behave.
That equation is the Henderson-Hasselbalch: pH = pKa + log([A⁻]/[HA]). It says the pH of a weak-acid solution depends not on absolute concentrations but on the ratio of the acid form to its conjugate base. The pKa is a property of the molecule; the ratio is something the environment can shift. From this one relationship falls everything from titration curves to the absorption of drugs across the gut wall to the way a kidney decides what to excrete. A buffer — a mixture of a weak acid and its conjugate base in roughly equal proportions — has the remarkable property of resisting pH change. Add H⁺, and the conjugate base absorbs it; add OH⁻, and the weak acid releases a proton to neutralize it. Buffers work best when the two forms are equal, which means pH ≈ pKa. The reason your blood sits in a band of 7.35–7.45, the reason a lake survives acid rain until it doesn't, the reason an aquarium can be kept alive at all: buffer chemistry, applied at every scale.
The most consequential present-tense story is ocean acidification. About a third of the CO₂ humans put into the atmosphere ends up dissolved in seawater, where it forms carbonic acid and shifts the carbonate buffer toward H⁺. Surface-ocean pH has fallen from preindustrial 8.2 to about 8.1 today — a 0.1-unit move that is, on a logarithmic scale, a roughly 30% increase in hydrogen-ion concentration, accomplished in two centuries against the backdrop of an ocean that had been chemically stable for tens of millions of years. Calcifying organisms — corals, oysters, the carbonate-shelled plankton that anchor marine food webs — are the most exposed: in a more acidic ocean, dissolved carbonate becomes scarce, and shells become harder to build and easier to dissolve. The Sørensen scale, designed for a brewery in 1909, now measures one of the largest chemistry experiments ever run on a planet.