P Ka Calculator

The acid dissociation constant Ka is the equilibrium constant for proton transfer from an acid to water, and because measured values span more than twenty orders of magnitude they are compressed into pKa, the negative base-ten logarithm. Each unit of pKa is a tenfold change in dissociation, so the scale behaves additively in energy terms. Representative values at 25 degrees C are hydrochloric acid near -6, trichloroacetic acid 0.66, formic acid 3.75, acetic acid 4.76, carbonic acid 6.35 for its first proton, ammonium 9.25, phenol 9.99 and water 15.7. For a conjugate pair the relation pKa + pKb = pKw holds, and since the ion product of water is 1.0 x 10^-14 at 25 degrees C that sum is taken as 14.00; a strong acid necessarily has a weak conjugate base. The pH of a weak monoprotic acid follows from setting the equilibrium expression against the initial concentration C, and when dissociation is small the square root approximation pH = -log10(sqrt(Ka x C)) applies. It is reliable only when the fraction ionised is under about five percent, which fails for dilute solutions or acids with pKa below roughly 2, where the full quadratic or a charge-balance solution is needed; it also breaks down below about 10^-6 M, where the autoionisation of water dominates. The Henderson-Hasselbalch form, pH = pKa + log10 of the ratio of conjugate base to acid, describes buffers and shows why pH equals pKa at the half-equivalence point of a titration. Two further limits deserve attention: pKw is temperature dependent, falling to about 13.26 at 50 degrees C so the 14 shortcut is wrong outside room temperature, and thermodynamic constants are defined in terms of activities, so measured apparent pKa values shift with ionic strength.

The framework begins with Svante Arrhenius, whose 1884 dissertation and 1887 papers proposed that electrolytes dissociate into ions in solution, work that earned the 1903 Nobel Prize in Chemistry. Wilhelm Ostwald applied the law of mass action to that picture in 1888, producing the dilution law that made the dissociation constant of a weak acid a measurable quantity and allowing acids to be ranked numerically for the first time. Soren Sorensen introduced the p notation in 1909 at the Carlsberg Laboratory in Copenhagen while studying the effect of acidity on enzymes in brewing, defining pH as the negative logarithm of hydrogen ion concentration; the same operator was later applied to Ka, Kb and Kw. Lawrence Henderson published the buffer relation in 1908 while investigating how blood maintains a nearly constant pH, and Karl Hasselbalch recast it in logarithmic form in 1917 using Sorensen's notation, giving the equation both men are now named for. The conceptual foundation was broadened in 1923 when Johannes Bronsted in Copenhagen and Thomas Lowry in England independently defined acids as proton donors and bases as proton acceptors, which explained conjugate pairs and made pKa a property of a species rather than of a substance class. Gilbert Lewis proposed his electron pair definition the same year. Compilations of measured pKa values, notably those assembled by Arthur Albert and E. P. Serjeant and by the IUPAC critical data series, turned the constant into a routine tool in pharmacology, where it governs the ionised fraction of a drug at physiological pH 7.4.