PKAs

pKa is the negative logarithm (base 10) of the acid dissociation constant Ka for the equilibrium HA ⇌ H+ + A−. It provides a quantitative measure of an acid’s strength in a given solvent, most commonly water at 25°C. A smaller pKa indicates a stronger acid, because the equilibrium favors dissociation more strongly.

In practical terms, pKa helps predict the protonation state of a species at a given pH. The Henderson-Hasselbalch equation, pH = pKa + log([A−]/[HA]), relates pH to the ratio of deprotonated and protonated forms. Polyprotic acids have several pKa values, each corresponding to a successive loss of a proton. For example, phosphoric acid has pKa values around 2.15, 7.20, and 12.35 for its three deprotonations.

Typical pKa values in water at 25°C span a wide range. Strong inorganic acids have negative pKa values in water, while very weak acids exceed 20. Common organic acids include acetic acid (pKa ~4.76), formic acid (pKa ~3.75), and ammonium (pKa ~9.25). Water itself has a pKa of about 15.7 for the reaction H2O ⇌ H+ + OH−. Carbonic acid shows pKa values of about 6.35 (first) and 10.33 (second). These values can shift with solvent, temperature, and ionic strength.

Notes: pKa is solvent-dependent and can differ markedly in nonaqueous media. Measurement methods include acid-base titration and spectroscopic approaches. In biochemistry and pharmacology, pKa values influence the charge state and behavior of molecules in biological systems.

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