Acid–base indicators Knowledge-Based Study

Study notes

1. Introduction to Acid–Base Indicators

   An acid–base indicator is a chemical substance that changes color reversibly depending on the pH of the solution to which it is added. These substances are widely used in chemical laboratories to estimate solution pH and to detect the equivalence point or endpoint in acid–base titrations.

   Chemically, most acid–base indicators are themselves complex weak organic acids or weak organic bases. The undissociated form of the molecule displays one characteristic color, whereas its conjugate base (or conjugate acid) form displays a distinctly different color due to differences in electronic conjugation and light absorption.

   Chemical Equilibrium of Weak Acid Indicators

   Consider a typical weak acid indicator represented by the general formula HIn. In an aqueous solution, it establishes the following dynamic equilibrium:

   $$\text{HIn}_{(aq)} + \text{H}_2\text{O}_{(l)} ightleftharpoons \text{H}_3\text{O}^+_{(aq)} + \text{In}^-_{(aq)}$$

   In this system, HIn represents the acid form (Color A), and $\text{In}^-$ represents the conjugate base form (Color B).

   Response to pH Changes (Le Chatelier's Principle)

   • Addition of acid (Low pH): Increasing $[\text{H}_3\text{O}^+]$ shifts the equilibrium position to the left according to Le Chatelier's principle. As a result, the concentration of $\text{HIn}$ dominates, and the solution exhibits Color A.
   • Addition of base (High pH): Added hydroxide ions ($\text{OH}^-$) react with and consume $\text{H}_3\text{O}^+$, lowering its concentration. The equilibrium shifts to the right, causing $[\text{In}^-]$ to dominate and the solution to exhibit Color B.

   The Color Change Interval

   The acid dissociation constant for an indicator, designated as $K_{\text{In}}$, is defined by the expression:

   $$K_{\text{In}} = \frac{[\text{H}_3\text{O}^+][\text{In}^-]}{[\text{HIn}]}$$

   Rearranging this relationship yields the Henderson–Hasselbalch form:

   $$\text{pH} = \text{p}K_{\text{In}} + \log\left(\frac{[\text{In}^-]}{[\text{HIn}]}\right)$$

   The human eye generally perceives the distinct color of one form only when its concentration is approximately ten times greater than that of the other form:

   • When $\frac{[\text{HIn}]}{[\text{In}^-]} \ge 10$, the solution appears entirely in Color A: $\text{pH} \le \text{p}K_{\text{In}} - 1$.
   • When $\frac{[\text{In}^-]}{[\text{HIn}]} \ge 10$, the solution appears entirely in Color B: $\text{pH} \ge \text{p}K_{\text{In}} + 1$.

   Therefore, the effective transition range of most simple indicators spans roughly two pH units centered around $\text{p}K_{\text{In}}$:

   $$\text{pH range} = \text{p}K_{\text{In}} \pm 1$$

   Common Indicators and Their Applications

   Different indicators possess different $\text{p}K_{\text{In}}$ values and therefore change color over different pH ranges. Some common examples include:

   • Methyl Orange: Red in acidic solutions (pH < 3.1) and yellow in basic solutions (pH > 4.4); transition range is 3.1–4.4.
   • Bromothymol Blue: Yellow in acidic solutions (pH < 6.0) and blue in basic solutions (pH > 7.6); transition range is 6.0–7.6.
   • Phenolphthalein: Colorless in acidic and neutral solutions (pH < 8.2) and pink/magenta in alkaline solutions (pH > 10.0); transition range is 8.2–10.0.

   Selecting an Indicator for a Titration

   For an accurate titration, an indicator must change color at or near the equivalence point of the reaction. This requires the indicator's transition range to overlap the steep, rapid pH jump on the titration curve:

   • Strong Acid – Strong Base: The equivalence point is at pH 7.0 with a wide vertical jump (pH 4 to 10), so bromothymol blue, phenolphthalein, or methyl red can be used.
   • Weak Acid – Strong Base: The equivalence point is basic (pH > 7), making phenolphthalein an ideal choice.
   • Strong Acid – Weak Base: The equivalence point is acidic (pH < 7), making methyl orange or methyl red appropriate.