By Claire Gillespie
Updated Aug 30, 2022
A titration is a precise analytical technique used to determine the concentration of an unknown solution by reacting it with a solution of known concentration. By measuring the volume of titrant required to reach the reaction’s endpoint, chemists can calculate the analyte’s molarity with confidence.
Place the analyte in an Erlenmeyer flask. Mount the titrant in a burette, ensuring the tap is sealed and the volume indicator is clear.
Slowly add the titrant to the analyte while gently swirling. Add a few drops of an appropriate indicator—phenolphthalein, for example, turns from pink in basic solutions to colorless in acidic conditions. The endpoint is marked by a permanent color change.
For a 1:1 stoichiometric relationship, the calculation is straightforward:
\[\text{M}_{\text{titrant}}\times V_{\text{titrant}} = \text{M}_{\text{analyte}}\times V_{\text{analyte}}\]
When the ratio differs, incorporate the stoichiometric factor. For example, to determine the concentration of NaOH in a 25 mL sample using 35 mL of 1.25 M HCl (1:1 ratio), compute: \[\text{M}_{\text{NaOH}} = \frac{1.25\,\text{M}\times35\,\text{mL}}{25\,\text{mL}} = 1.75\,\text{M}\]
Titrations are foundational in laboratories worldwide, enabling accurate quantification of acids, bases, and other reactive species. They are critical in pharmaceuticals, environmental testing, and quality control, ensuring products meet stringent safety and efficacy standards.
Use the 1:1 molarity equation for most acid–base titrations. Adjust for stoichiometry when dealing with polyprotic acids or complex ions.
