Titration is a technique used in quantitative chemical analysis to determine the unknown concentration of a substance (the analyte) by reacting it with a solution of known concentration (the titrant). The titrant is added slowly to the analyte until the reaction is complete. An indicator is a substance added to the analyte solution to provide a clear, visual signal of this completion. The indicator marks the end of the chemical reaction with an abrupt color change, allowing a chemist to precisely measure the volume of titrant required.
How Indicators Chemically Signal Change
The ability of an indicator to change color is rooted in its chemical structure and a reversible equilibrium reaction sensitive to the solution’s environment. Most acid-base indicators function as weak acids (HIn) that exist in balance with their conjugate base (In-). The protonated (HIn) and deprotonated (In-) forms have distinctly different molecular structures, causing them to absorb and reflect different wavelengths of light, resulting in two separate colors.
In a highly acidic solution, high hydrogen ion concentration forces the equilibrium to favor the HIn form. As the titrant is added, it consumes hydrogen ions, causing the pH to rise. This shift causes the equilibrium to move away from the HIn form and toward the In- form, resulting in a sudden color transition.
This color change occurs over a narrow range of pH values, known as the transition range. The transition is sharp because the indicator is designed to switch forms rapidly when the concentration of the titrant becomes sufficient.
Matching the Indicator to the Titration
The success of a titration depends on matching the indicator’s color-change range to the pH of the equivalence point. The equivalence point is the theoretical moment when the exact stoichiometric amount of titrant has been added. The endpoint is the experimental point where the indicator physically changes color. For accurate results, the endpoint must occur as close as possible to the equivalence point to minimize error.
Indicator selection is guided by the titration curve, a graph plotting the solution’s pH against the volume of titrant added. This curve reveals a steep vertical region where the pH changes dramatically with the addition of a tiny volume of titrant. The equivalence point sits within this steep region. A suitable indicator must have its pH transition range, typically about two pH units wide, spanning the pH of this equivalence point.
For example, a strong acid/strong base titration has a neutral equivalence point (pH 7.0), making bromothymol blue appropriate. A weak acid/strong base titration results in a basic equivalence point (around pH 8 to 9), requiring phenolphthalein. A weak base/strong acid titration has an acidic equivalence point (around pH 3 to 5), requiring methyl orange. Selecting the correct indicator ensures the visible endpoint coincides with the theoretical equivalence point.
Main Categories of Titration Indicators
While acid-base indicators are the most common, indicators are used in several other types of titrations, each relying on a different chemical principle for the visual signal. Redox indicators, for example, do not respond to pH changes but rather to a shift in the solution’s reduction potential. These molecules change color when they are oxidized or reduced by the titrant, signaling the completion of an electron transfer reaction. Starch solution, which forms a deep blue complex with iodine, is used in certain iodine-based redox titrations.
Precipitation titrations employ indicators to mark the formation of an insoluble solid between the analyte and the titrant. The indicator reacts with the first slight excess of titrant after the analyte is consumed, forming a colored precipitate or soluble complex. Complexometric titrations, which involve the formation of a stable, soluble complex, often use metal ion indicators that change color when the metal ion concentration drops sharply at the equivalence point.