Why Do Indicators Change Color?
Analogy- You're conducting a titration, carefully adding a base to an acidic solution drop by drop.
- Suddenly, the solution shifts from clear to pink: an unmistakable signal that the reaction has reached its endpoint.
- But have you ever wondered why this dramatic color change occurs? What’s happening at the molecular level to produce such a visible transformation?
- The answer lies in the chemistry of acid–base indicators, compounds specifically designed to reveal subtle changes in pH.
Acid–Base Indicators: Weak Acids with a Colorful Twist
Definition
Acid–base indicators
Acid–base indicators are weak acids that exist in equilibrium between their protonated form (HInd) and deprotonated form (Ind⁻).
- It can be expressed by: $$
\text{HInd (aq)} \rightleftharpoons \text{H}^+ \text{(aq)} + \text{Ind}^- \text{(aq)}
$$ - What makes these compounds unique is that HInd and Ind⁻ have different colors. For instance:
- HInd might be red.
- Ind⁻ might be blue.
The observed color of the solution depends on the relative concentrations of HInd and Ind⁻, which are influenced by the solution’s pH.
- At low pH (acidic conditions), the equilibrium shifts toward HInd, coloring the solution red.
- At high pH (basic conditions), Ind⁻ dominates, turning the solution blue.
Indicators are not "magic"; they follow the same equilibrium principles as other weak acids!
Equilibrium and Observed Color
- To pinpoint when an indicator changes color, we use its equilibrium expression: $$
K_a = \frac{[\text{H}^+][\text{Ind}^-]}{[\text{HInd}]}
$$ - Rewriting this in terms of pH and $pK_a$: $$
\text{pH} = \text{pKa} + \log \left( \frac{[\text{Ind}^-]}{[\text{HInd}]} \right)
$$
Key Observations:
- At low pH: $[\text{H}^+]$ is high, so the equilibrium shifts left, favoring HInd. The solution takes on HInd’s color.
