Titrimetry

Titrimetry, or titration, is a quantitative chemical analysis in which a solution of known concentration, the titrant, reacts with a solution of unknown concentration, the analyte. The number of milliliters of titrant required to reach the reaction’s completion is used, with the reaction’s stoichiometry, to calculate the amount of analyte present.

The point at which the reaction is complete is called the end point; the equivalence point is the theoretical point of exact stoichiometric balance. In practice end points are detected by indicators that change color, by pH or voltage measurements with a meter, or by conducting the titration and observing a change in analytical signal.

Common types include acid-base titrations, redox titrations, complexometric titrations, and precipitation titrations. In acid-base titrations, indicators such as phenolphthalein or methyl orange are used, or a pH meter can determine the endpoint. Redox titrations use oxidizing or reducing agents, for example permanganate or iodine solutions. Complexometric titrations, notably with EDTA, determine metal ions by forming stable complexes. Precipitation titrations involve a reaction that forms a precipitate, often detected by turbidity or by cessation of reaction with a silver salt. Non-aqueous titrations use solvents other than water and can extend titrimetry to non-aqueous samples.

Standardization of the titrant against a primary standard ensures accuracy; calculations follow from the volume at endpoint and the stoichiometry. Modern titration often employs automatic titrators and potentiometric detection, which can improve precision and reduce operator bias.

Applications include determining acids and bases in pharmaceuticals, environmental analysis, water hardness, food chemistry, and clinical chemistry. Historically, titrimetry developed in the 19th century with methods introduced by Karl Friedrich Mohr and others and remains a fundamental technique in analytical chemistry.