91Ó°ÊÓ

How does an acid-base indicator work?

A student carried out two titrations using a \(\mathrm{NaOH}\) solution of unknown concentration in the buret. In one titration she weighed out \(0.2458 \mathrm{~g}\) of KHP (see Section 4.7 ) and transferred it to an Erlenmeyer flask. She then added \(20.00 \mathrm{~mL}\) of distilled water to dissolve the acid. In the other titration she weighed out \(0.2507 \mathrm{~g}\) of KHP but added \(40.00 \mathrm{~mL}\) of distilled water to dissolve the acid. Assuming no experimental error, would she obtain the same result for the concentration of the \(\mathrm{NaOH}\) solution?

Would the volume of a \(0.10 M \mathrm{NaOH}\) solution needed to titrate \(25.0 \mathrm{~mL}\) of a \(0.10 \mathrm{M} \mathrm{HNO}_{2}\) (a weak acid) solution be different from that needed to titrate \(25.0 \mathrm{~mL}\) of a \(0.10 \mathrm{M} \mathrm{HCl}\) (a strong acid) solution?

A quantity of \(18.68 \mathrm{~mL}\) of a KOH solution is needed to neutralize \(0.4218 \mathrm{~g}\) of \(\mathrm{KHP}\). What is the concentration (in molarity) of the KOH solution?

Calculate the concentration (in molarity) of a \(\mathrm{NaOH}\) solution if \(25.0 \mathrm{~mL}\) of the solution are needed to neutralize \(17.4 \mathrm{~mL}\) of a \(0.312 \mathrm{M} \mathrm{HCl}\) solution.

Calculate the volume in milliliters of a \(1.420 \mathrm{M}\) \(\mathrm{NaOH}\) solution required to titrate the following solutions. (a) \(25.00 \mathrm{~mL}\) of a \(2.430 \mathrm{M} \mathrm{HCl}\) solution (b) \(25.00 \mathrm{~mL}\) of a \(4.500 \mathrm{M} \mathrm{H}_{2} \mathrm{SO}_{4}\) solution (c) \(25.00 \mathrm{~mL}\) of a \(1.500 \mathrm{M} \mathrm{H}_{3} \mathrm{PO}_{4}\) solution

What volume of a \(0.500 M\) HCl solution is needed to neutralize each of the following? (a) \(10.0 \mathrm{~mL}\) of a \(0.300 \mathrm{M} \mathrm{NaOH}\) solution (b) \(10.0 \mathrm{~mL}\) of a \(0.200 \mathrm{M} \mathrm{Ba}(\mathrm{OH})_{2}\) solution

What are the similarities and differences between acid-base titrations and redox titrations?

Explain why potassium permanganate \(\left(\mathrm{KMnO}_{4}\right)\) and potassium dichromate \(\left(\mathrm{K}_{2} \mathrm{Cr}_{2} \mathrm{O}_{7}\right)\) can serve as internal indicators in redox titrations.

Iron(II) can be oxidized by an acidic \(\mathrm{K}_{2} \mathrm{Cr}_{2} \mathrm{O}_{7}\) solution according to the net ionic equation \(\mathrm{Cr}_{2} \mathrm{O}_{7}^{2-}+6 \mathrm{Fe}^{2+}+14 \mathrm{H}^{+} \longrightarrow\) \(2 \mathrm{Cr}^{3+}+6 \mathrm{Fe}^{3+}+7 \mathrm{H}_{2} \mathrm{O}\) If it takes \(26.0 \mathrm{~mL}\) of \(0.0250 \mathrm{M} \mathrm{K}_{2} \mathrm{Cr}_{2} \mathrm{O}_{7}\) to titrate \(25.0 \mathrm{~mL}\) of a solution containing \(\mathrm{Fe}^{2+},\) what is the molar concentration of \(\mathrm{Fe}^{2+} ?\)