Question

Acid-base indicators mark the end point of titrations by "magically" turning a different color. Explain the "magic" behind acid-base indicators.

Short answer

The "magic" behind acid-base indicators is due to the equilibrium between their protonated (HIn) and deprotonated (In-) forms, each displaying a distinct color. As the pH of the solution changes during titration, the ratio of HIn to In- also changes, causing a visible color change when their concentrations are equal. The exact pH at which this occurs depends on the indicator's pKa value, following the Henderson-Hasselbalch equation. Factors affecting the color change include the specific indicator chosen, its concentration, and the pH range of the titration.

Step-by-step solution

1. Acid-base indicators and their role in titration

Acid-base indicators are weak acids or bases that display different colors in their ionized and unionized forms. In titrations, they are used to determine the end point – the point at which the reaction between the analyte and the titrant is complete. At the end point, the indicator changes its color, signaling the completion of the reaction.

2. How color change in indicators occur

Acid-base indicators are composed of molecules that exist in equilibrium between their protonated (HIn) and deprotonated (In-) forms. Each form has a distinct color. When the pH of the solution changes during titration, the ratio of HIn to In- also changes, which eventually causes a visible color change. The equation representing this equilibrium is: \[HIn(aq) \rightleftharpoons H^+(aq) + In^-(aq)\] The color change occurs when the concentrations of the protonated HIn and deprotonated In- forms are equal. The exact pH at which this occurs depends on the indicator's pKa value, as given by the Henderson-Hasselbalch equation: \[\text{pH} = \text{p}K_a + \log_{10}\Big(\frac{[\text{In}^-]}{[\text{HIn}]}\Big)\]

3. Factors affecting the color change

The color change of an acid-base indicator depends on: 1. The specific indicator chosen: Different indicators have different pKa values and distinctive color changes that occur at different pH levels. 2. The concentration of the indicator: Higher concentrations can make the color change more intense and easier to see. 3. The pH range of the titration: An appropriate indicator should be chosen based on the pH range of the titration. For instance, if the titration is expected to have a pH range of 4-6, an indicator with a pKa value close to 5 should be chosen. In summary, the "magic" behind acid-base indicators lies in the equilibrium between their protonated and deprotonated forms along with their distinctive colors in different forms and pH-dependent color changes, which allow them to signal the end point of a titration.