pH

By 1909, chemistry faced an irritating problem: the concentration of hydrogen ions in solution—critical for understanding everything from enzyme activity to beer quality—required writing numbers like 0.0000001 molar. Svante Arrhenius's 1887 ion theory had established that acids dissociate into hydrogen ions in water, giving chemists the conceptual framework to understand acidity quantitatively. But the numbers themselves were unwieldy.

The adjacent possible was converging. Electrolytic dissociation theory provided the mechanism. Electrochemical measurement techniques offered ways to measure hydrogen ion concentration via electrical potential. Logarithmic scales were well-established mathematical tools. All that was needed was someone facing a practical problem urgent enough to demand standardization.

That person was Søren Sørensen, head of chemistry at the Carlsberg Laboratory in Copenhagen. Carlsberg wasn't just a brewery—it was a research institution where scientific rigor met industrial necessity. Sørensen was studying how ion concentrations affected proteins, particularly brewing enzymes. Each enzyme had a narrow optimal range, but expressing this required saying things like "the enzyme works best at hydrogen ion concentration of 7.2×10⁻⁵ molar."

In 1909, Sørensen introduced pH: the negative logarithm of hydrogen ion concentration. The "p" likely stood for "power" (potenz in German). Suddenly, that clunky number became pH 4.14. Pure water's 1×10⁻⁷ became the elegant pH 7. The scale compressed a million-fold range into a manageable 0-14 span.

The cascade was immediate. By the 1920s, pH became standard in physiology. The concept's full potential unlocked in 1934 when Arnold Beckman invented the pH meter—portable, robust, usable by non-specialists. pH measurement became routine, enabling real-time monitoring in fermentation, drug synthesis, environmental testing, and clinical diagnostics.