Titration

Titration—also known as titrimetry or volumetric analysis—is a foundational laboratory technique in analytical chemistry used to determine the concentration of an analyte within a solution. The method involves reacting a solution of unknown concentration with a reagent of precisely known concentration, known as the titrant. By measuring the volume of titrant required to complete the reaction, chemists can calculate the concentration of the analyte with high accuracy. Titrations are integral to quantitative analysis across chemical, biological, industrial, pharmaceutical, and environmental laboratories.

Historical Development and Etymology

The terminology and methodology of titration emerged gradually through developments in eighteenth- and nineteenth-century French chemistry. The word derives from the French titrer, originally meaning the proportion of precious metal in coins or metalwork—a measure of fineness that evolved into the concept of concentration. By 1828, Joseph Louis Gay-Lussac had transformed titre into a verb meaning to determine concentration, marking an important lexical shift towards modern analytical usage.
Volumetric techniques were first systematised in late eighteenth-century France. François-Antoine-Henri Descroizilles produced an early burette resembling a graduated cylinder in 1791, marking one of the earliest purpose-built titration instruments. Gay-Lussac later refined this tool by adding a side arm and introduced the terms pipette and burette in his 1824 work on indigo standardisation.
The true modern burette emerged in 1845 through Étienne-Ossian Henry’s design, which incorporated features recognisable in contemporary glassware. Karl Friedrich Mohr’s mid-nineteenth-century improvements—including simplified construction, practical stopcocks, and standardised procedures—popularised titration and laid the foundation for analytical teaching through his 1855 textbook on titration methodology.

General Procedure

A typical titration involves placing a carefully measured quantity of analyte into a beaker or Erlenmeyer flask, often with a few drops of an appropriate indicator. A calibrated burette containing the titrant is positioned above the vessel. The titrant is added gradually while the solution is mixed, allowing the reagent to react with the analyte.
The endpoint of the titration is detected when the indicator undergoes a permanent colour change, signifying that stoichiometric equivalence—known as the equivalence point—has been reached or closely approached. The precision of the result depends on careful measurement of the titration volume and appropriate selection of indicator. In highly sensitive analyses, instrumental detection such as potentiometry or conductometry replaces visual endpoints.

Sample Preparation and Experimental Considerations

Successful titration requires the analyte and titrant to be in solution. Solid analytes are often dissolved in water, though non-aqueous solvents such as acetic acid or ethanol are used when water interferes with the reaction. Dilution is frequently employed to improve accuracy and control reaction kinetics.
In many complexometric or redox titrations, maintaining a constant pH is essential; buffer solutions are therefore added to stabilise conditions. When samples contain multiple reactive species, a masking agent may be used to suppress undesired reactions and isolate the analyte’s behaviour.
Some redox reactions, such as the oxidation of oxalate, require elevated temperatures to achieve reasonable reaction rates. In such cases the titration proceeds while the solution remains hot, with careful control to prevent decomposition or side reactions.

Titration Curves and Their Interpretation

A titration curve plots the progression of the titration through a graph in which the x-axis represents the volume of titrant added and the y-axis represents a measurable property of the solution—commonly pH in acid–base systems. The curve provides insight into reaction behaviour, buffering regions, and the equivalence point.
Strong acid–strong base titrations produce smooth, sharply rising curves near the equivalence point, enabling comfortable use of various indicators such as phenolphthalein, bromothymol blue, or litmus. When weak acids or weak bases are involved, the curves become less steep, and the indicator must be carefully matched to the pH range around the equivalence point.
For weak acid–strong base titrations, the solution becomes basic at the equivalence point, favouring indicators like phenolphthalein. In weak base–strong acid titrations, the equivalence point is acidic, and indicators such as methyl orange or bromothymol blue are more suitable. In titrations involving both weak acid and weak base, the curve may be too irregular for visual detection, necessitating instrumental methods such as pH metres.
Titration curves typically follow a sigmoid shape, reflecting the gradual consumption of the analyte and the rapid change near equivalence.

Types of Titrations

Titration methods are diverse and tailored to specific analytical goals. The most common include:

• Acid–base titrations, based on neutralisation reactions and monitored by indicators or pH measurement. They require careful consideration of acid and base strengths, buffer effects, and the role of conjugate species.
• Redox titrations, involving electron-transfer reactions. These often use oxidising agents such as potassium permanganate or reducing agents such as sodium thiosulfate, and may rely on self-indicating reagents or electrochemical detection.

Beyond these principal methods, numerous specialised titrations exist, including precipitation titrations, complexometric titrations, and non-aqueous titrations for strong bases insoluble in water. Advanced organometallic titrations, for example, use anhydrous solvents and weakly acidic indicators to avoid interference from aqueous protonation.

Quantitative Treatment

Titration calculations depend on stoichiometry and equilibrium principles. Before titrant addition, the analyte concentration follows standard weak acid or base dissociation relationships. At equivalence, the pH is determined by hydrolysis of the reaction products. Between these extremes, the Henderson–Hasselbalch equation describes buffer behaviour. Accurate calculation may require simultaneous treatment of acid dissociation, base hydrolysis, and water self-ionisation through a system of equilibrium equations.