Indicators
Core Concept
Definition: Acid-base indicators are substances that change color depending on the pH of the solution.
Key Property: They exist in equilibrium between two forms:
Acidic form (HIn): Dominates in acidic solutions.
Basic form (In⁻): Dominates in basic solutions.
Indicators are weak acids or bases that change color at specific pH ranges.
The choice of an indicator depends on the reaction's equivalence point.
Natural and synthetic indicators can be used for titrations and pH testing.
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How Do Indicators Work?
Indicators are weak acids or bases that dissociate in solution according to the following equilibrium:
$$\text{HIn} \rightleftharpoons \text{H}^+ + \text{In}^-$$
Color Change
In acidic solutions ($\text{pH} < \text{p}K_a$): The equilibrium shifts left (Le Châtelier's principle). The acidic form ($\text{HIn}$) dominates, giving one color.
In basic solutions ($\text{pH} > \text{p}K_a$): The equilibrium shifts right. The basic form ($\text{In}^-$) dominates, giving a different color.
Key Terms
pH Range: The range of pH values over which the indicator changes color.
Transition Point: The pH at which the indicator is halfway between its acidic and basic forms (pH=pKa\text{pH} = \text{p}K_apH=pKa).
| Indicator | Color in Acid | Color in Base | Transition Range (pH) |
|---|---|---|---|
| Litmus | Red | Blue | 4.5–8.3 |
| Methyl orange | Red | Yellow | 3.1–4.4 |
| Bromothymol blue | Yellow | Blue | 6.0–7.6 |
| Phenolphthalein | Colorless | Pink | 8.2–10.0 |
| Universal indicator | Red | Violet | 4.0–10.0 (varies) |
Choosing the Right Indicator
Equivalence Point: The indicator must change color close to the pH of the reaction's equivalence point.
For strong acid-strong base titrations: Use indicators like phenolphthalein or bromothymol blue (pH ≈ 7).
For weak acid-strong base titrations: Use phenolphthalein (pH > 7 at equivalence).
For strong acid-weak base titrations: Use methyl orange (pH < 7 at equivalence).
Common Acid-Base Indicators