Selection of an Indicator for Titrations
Understanding Key Terms: Equivalence Point vs. End Point
Before selecting an indicator, it's essential to distinguish two key terms:
- Equivalence Point: The point in a titration where the acid and base have reacted in exact stoichiometric proportions.
- End Point: The moment when the indicator changes color, signaling the titration's completion. Ideally, the end point should closely align with the equivalence point.
$$
\text{HCl(aq) + NaOH(aq) → NaCl(aq) + H}_2\text{O(l)}
$$
The equivalence point occurs when 1 mole of HCl reacts with 1 mole of NaOH.
The Role of pH at the Equivalence Point
The pH at the equivalence point depends on the type of acid and base involved:
- Strong Acid + Strong Base: $pH = 7$ (neutral solution of salt and water).
- Weak Acid + Strong Base: $pH > 7$ (basic solution due to conjugate base hydrolysis).
- Strong Acid + Weak Base: $pH < 7$ (acidic solution due to conjugate acid hydrolysis).
- Weak Acid + Weak Base: pH depends on the relative strengths of the acid and base, often near neutral.
Determine the pH at the equivalence point to guide your choice of indicator.
How Indicators Work
AnalogyThink of an indicator as a chemical "mood ring." It changes color depending on the pH of the solution.
- Indicators are weak acids or bases that exist in two forms: protonated and deprotonated.
- Each form has a distinct color. $$
\text{HInd(aq)} ⇌ \text{H}^+\text{(aq)} + \text{Ind}^-\text{(aq)}
$$ - The color change occurs over a pH range of approximately pKa ± 1.
