Acids and Bases
Acid-Base Indicators
Acid-base indicators are substances that change color depending on the pH of the solution they are in. This color change is due to a change in the indicator’s chemical structure, often involving protonation or deprotonation.
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Mechanism: Indicators are typically weak acids or bases. The equilibrium between the protonated and deprotonated forms determines the color observed. Consider a weak acid indicator, HIn:
$ HIn \rightleftharpoons H^+ + In^- $
The equilibrium constant for this reaction is the acid dissociation constant, $ K_a $ .
$$ K_a = \frac{[H^+]][In^-]]}{[HIn]]} $$
The ratio of $ [In^-]] $ to $ [HIn]] $ determines the color observed. At low pH (high $ [H^+]]) $ , the equilibrium shifts to the left, favoring the protonated form ( $ HIn $ ), resulting in one color. At high pH (low $ [H^+]]) $ , the equilibrium shifts to the right, favoring the deprotonated form ( $ In^- $ ), resulting in a different color. The pKa of the indicator determines the pH range over which the color change occurs. The transition range is typically approximately one pH unit around the pKa value.
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Examples:
- Phenolphthalein: Colorless in acidic solution, pink in basic solution. Transition range: ~8.2-10.0
- Methyl orange: Red in acidic solution, yellow in basic solution. Transition range: ~3.1-4.4
- Bromothymol blue: Yellow in acidic solution, blue in basic solution. Transition range: ~6.0-7.6
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Choosing an indicator: The appropriate indicator for a titration depends on the pH at the equivalence point. The indicator’s transition range should encompass the equivalence point pH. Titration Curves
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Limitations: Indicators provide an approximate measure of pH. They are not precise instruments for determining pH. The exact pH at which the color change occurs can be influenced by temperature and ionic strength.