By Brooke Yool • Updated August 30, 2022
In a titration, you determine the concentration of an unknown solution by adding a reagent of known concentration. The endpoint occurs when the stoichiometric ratio of the two reactants is exactly met—no excess of either. The volume of titrant used then allows you to calculate the unknown concentration.
Because most reactions are not visually distinct at the endpoint, we add a pH‑sensitive or complex‑forming indicator. When the reaction reaches equivalence, the indicator’s molecular structure changes, producing a visible color shift that signals completion.
Acid–base titrations demand an indicator whose color transition aligns with the expected endpoint pH. By sketching a titration curve (pH vs. titrant volume), you can select an indicator that changes color at the steepest slope—typically within 0.5 pH units of the equivalence point.
Anthocyanin pigments in plants serve as classic natural indicators. For example, a red‑cabbage extract turns bluish‑green in a base and pink in an acid. Flower petals can also shift color with soil pH, illustrating the same principle.
Temperature and pressure variations can alter an indicator’s response, and pH indicators provide only approximate values. For precise measurements, use a calibrated pH meter instead of relying solely on color change.
