Acid–base indicators Knowledge-Based Download PDF

Questions

1. 1.

   Why is only a very small quantity (a few drops) of an acid–base indicator typically added to a titration flask?

   • A) Because indicators are weak acids or bases that react with the titrant, and excessive amounts would alter the required titration volume.
   • B) Because large quantities of indicator would permanently shift the stoichiometric equivalence point of any titration to exactly pH 7.0.
   • C) Because excessive indicator suppresses the ionization of the titrant by completely absorbing visible light.
   • D) Because high indicator concentrations precipitate out of aqueous solution and prevent any visual color change from occurring.
2. 2.

   A typical acid–base indicator changes color over an interval of approximately two pH units centered around its $\text{p}K_{\text{In}}$. Which of the following explains why this visual transition spans approximately two pH units?

   • A) The indicator decomposes irreversibly when the hydronium ion concentration changes by more than one order of magnitude.
   • B) The autoionization of water restricts the protonation–deprotonation equilibrium of all weak acids to a two-unit span.
   • C) An observer generally requires an approximate tenfold excess of one colored conjugate form relative to the other to perceive a definitive color change.
   • D) The indicator undergoes two successive proton transfer reactions that each shift the solution by one pH unit.
3. 3.

   A student titrates a sample of acetic acid, a weak acid, with sodium hydroxide, a strong base, resulting in an equivalence point at approximately $\text{pH} = 8.7$. Which indicator is most suitable for detecting the equivalence point of this titration?

   • A) Methyl orange (transition range $\text{pH} \approx 3.1 - 4.4$)
   • B) Bromothymol blue (transition range $\text{pH} \approx 6.0 - 7.6$)
   • C) Phenolphthalein (transition range $\text{pH} \approx 8.2 - 10.0$)
   • D) Methyl red (transition range $\text{pH} \approx 4.4 - 6.2$)
4. 4.

   An acid–base indicator, $HIn$, has an ionization constant of $K_{\text{In}} = 1.0 \times 10^{-5}$. Its acidic form ($HIn$) is yellow and its conjugate base form ($\text{In}^-$) is blue. What color will the indicator display in a solution buffered at pH 7.0, and which species predominates?

   • A) Yellow, because pH is higher than $\text{p}K_{\text{In}}$ and $HIn$ predominates
   • B) Green, because pH equals $\text{p}K_{\text{In}}$ and both forms are present in equal amounts
   • C) Yellow, because the indicator neutralizes the solution and converts entirely into its acid form
   • D) Blue, because pH is higher than $\text{p}K_{\text{In}}$ and $\text{In}^-$ predominates
5. 5.

   Which statement correctly describes the chemical nature and operating principle of a common acid–base indicator in aqueous solution?

   • A) It is a strong acid that neutralizes excess hydroxide ions to shift the equivalence point of a titration.
   • B) It is a weak acid or weak base whose protonated and deprotonated forms display different colors.
   • C) It is an inert dye whose color depends solely on the temperature and ionic strength of the solvent.
   • D) It is a concentrated buffer solution that resists any change in pH across an acid–base titration.