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 lies in the chemical equilibrium between two different forms of the indicator molecule, which have different colors. Indicators exist as weak acids (HIn) or weak bases (In^-) and can lose or gain a proton based on the pH of the solution. The equilibrium can be represented as: \( HIn \rightleftharpoons H^+ + In^- \). At low pH, the equilibrium shifts to the left, displaying the color associated with HIn; while at high pH, the equilibrium shifts to the right, displaying the color associated with In^-. The pH at which the color change occurs corresponds to the pKa of the indicator.

Step-by-step solution

Introduction to acid-base indicators

An acid-base indicator is a substance that can be added to a solution to help determine the end point of a titration, usually an acid-base titration. In a titration, it is important to know when the reaction has reached its equivalence point or the point at which the amount of acid equals the amount of base. Acid-base indicators accomplish this by changing color in response to the pH of the solution, indicating if the solution is acidic or basic.

Color change in acid-base indicators

The color change observed in acid-base indicators is due to the presence of two different chemical forms of the indicator molecule that are interconvertible in response to the pH of the solution. These two forms of the indicator have different colors. When the pH of the solution changes, the proportion of these two forms shifts, leading to a different color being displayed. The pH at which the color change occurs is called the transition pH or pKa of the indicator.

Chemical equilibrium of acid-base indicators

The chemistry behind the color change of acid-base indicators lies in the equilibrium that exists between the two forms of the indicator. The indicators exist as weak acids (abbreviated as HIn) or weak bases (In^-) which can lose or gain a proton, respectively. The chemical equilibrium can be represented as: \( HIn \rightleftharpoons H^+ + In^- \) At low pH (acidic solution), the equilibrium shifts to the left, resulting in a higher concentration of HIn and displaying the color associated with this form of the indicator. At high pH (basic solution), the equilibrium shifts to the right, leading to a higher concentration of In^-, and the color associated with this form is displayed. The pKa of the indicator determines the pH at which the color change occurs. When the pH of the solution is equal to the pKa of the indicator, equal amounts of both forms (HIn and In^-) are present in the solution. By understanding the chemical equilibrium of acid-base indicators and their color change based on the pH of the solution, we can explain the "magic" behind the color-changing properties of these valuable tools in titrations.