Titrations Based on Acid–Base Reactions
The earliest acid–base titrations involved the determination of the acidity or alka- linity of solutions, and the purity of carbonates and alkaline earth oxides. Before 1800, acid–base titrations were conducted using H2SO4, HCl, and HNO3 as acidic titrants, and K2CO3 and Na2CO3 as basic titrants. End points were determined using visual indicators such as litmus, which is red in acidic solutions and blue in basic solutions, or by observing the cessation of CO2 effervescence when neutraliz- ing CO32–. The accuracy of an acid–base titration was limited by the usefulness of the indicator and by the lack of a strong base titrant for the analysis of weak acids.
The utility of acid–base titrimetry improved when NaOH was first introduced as a strong base titrant in 1846. In addition, progress in synthesizing organic dyes led to the development of many new indicators. Phenolphthalein was first synthe- sized by Bayer in 1871 and used as a visual indicator for acid–base titrations in 1877. Other indicators, such as methyl orange, soon followed. Despite the increas- ing availability of indicators, the absence of a theory of acid–base reactivity made se- lecting a proper indicator difficult.
Developments in equilibrium theory in the late nineteenth century led to sig- nificant improvements in the theoretical understanding of acid–base chemistry and, in turn, of acid–base titrimetry.
Sørenson’s establishment of the pH scale in 1909 provided a rigorous means for comparing visual indicators. The determination of acid–base dissociation constants made the calculation of theoretical titration curves possible, as outlined by Bjerrum in 1914. For the first time a rational method ex- isted for selecting visual indicators, establishing acid–base titrimetry as a useful al- ternative to gravimetry.