Why Do Indicators Change Color?
- 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
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.
- At high pH: $[\text{H}^+]$ is low, so the equilibrium shifts right, favoring Ind⁻. The solution takes on Ind⁻’s color.
- At pH = pKa: $[\text{HInd}] = [\text{Ind}^-]$, and the color is a mix of the two forms.
Imagine using bromothymol blue, an indicator with a pKa of 7.0.
- In acidic solutions (pH< 6.0), the solution appears yellow because HInd dominates.
- In basic solutions (pH >7.6), it turns blue as Ind⁻ takes over. At neutral pH (around 7.0), the solution is green—a mix of yellow and blue.
Don’t confuse the endpoint of an indicator with the equivalence point of a titration! They only match if the indicator’s pKa aligns with the titration’s equivalence point.
Examples of Indicators
Different indicators are suited to different pH ranges, depending on their pKa values. Here are some common examples:
| Indicator | $\text{pK}_a$ | pH Transition Range | Color Change |
|---|---|---|---|
| Methyl orange | 3.7 | 3.1-4.4 | $\text{Red} \rightarrow \text{Orange} \rightarrow \text{Yellow}$ |
| Bromothymol blue | 7.0 | 6.0-7.6 | $\text{Yellow} \rightarrow \text{Green} \rightarrow \text{Blue} $ |
| Phenolphthalein | 9.6 | 8.3-10.0 | $\text{Colorless} \rightarrow \text{Pink}$ |
When choosing an indicator for a titration, match its pH transition range to the pH at the equivalence point.
Universal Indicator: A Symphony of Colors
- Unlike single indicators, the universal indicator is a mixture of several indicators, each with a different pKa.
- This combination allows it to display a continuous range of colors across the entire pH spectrum, from red (pH 1) to violet (pH 14).
The universal indicator is not a single chemical: it’s a carefully designed blend of multiple indicators.
- What is the relationship between an indicator’s pKa and the pH at which it changes color?
- Why do different indicators have different pH ranges?
- How would you choose the best indicator for a titration involving a weak acid and a strong base?
- Could a single indicator ever provide precise results across the entire pH scale? Why or why not?
