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Chapter 15: Problem 18

Which of the following can be classified as buffer solutions? a. \(0.25 \mathrm{M} \mathrm{HBr}+0.25 \mathrm{M} \mathrm{HOBr}\) b. \(0.15 \mathrm{M} \mathrm{HClO}_{4}+0.20 \mathrm{M} \mathrm{RbOH}\) c. \(0.50 \mathrm{M} \mathrm{HOCl}+0.35 \mathrm{M} \mathrm{KOCl}\) d. \(0.70 \mathrm{MKOH}+0.70 \mathrm{M} \mathrm{HONH}_{2}\) e. \(0.85 \mathrm{M} \mathrm{H}_{2} \mathrm{NNH}_{2}+0.60 \mathrm{M} \mathrm{H}_{2} \mathrm{NNH}_{3} \mathrm{NO}_{3}\)

Short Answer

Expert verified
Option c: \(0.50\,\mathrm{M}\,\mathrm{HOCl} + 0.35\,\mathrm{M}\,\mathrm{KOCl}\) and option e: \(0.85\,\mathrm{M}\,\mathrm{H}_2\,\mathrm{NNH}_2 + 0.60\,\mathrm{M}\, \mathrm{H}_2\,\mathrm{NNH}_3\,\mathrm{NO}_3\) can be classified as buffer solutions.

Step by step solution

01

Option a: \(0.25\,\mathrm{M}\,\mathrm{HBr} + 0.25\,\mathrm{M}\,\mathrm{HOBr}\)

Here, HBr is a strong acid, and HOBr is a weak acid. Since the solution contains a strong acid, it cannot be a buffer solution, as the presence of strong acids makes the solution unable to resist pH changes.
02

Option b: \(0.15\,\mathrm{M}\,\mathrm{HClO}_4 + 0.20\,\mathrm{M}\,\mathrm{RbOH}\)

HClO4 is a strong acid, and RbOH is a strong base. The presence of both a strong acid and strong base will neutralize the solution, but it will not create a buffer solution.
03

Option c: \(0.50\,\mathrm{M}\,\mathrm{HOCl} + 0.35\,\mathrm{M}\,\mathrm{KOCl}\)

This solution contains HOCl (a weak acid) and its conjugate base OCl- (from KOCl). Both the weak acid and its conjugate base are present in significant amounts in the solution, making this a buffer solution.
04

Option d: \(0.70\,\mathrm{M}\,\mathrm{KOH} + 0.70\,\mathrm{M}\,\mathrm{HONH}_2\)

In this case, KOH is a strong base, and HONH2 is a weak acid. As we have a strong base present, the solution will not be able to resist changes in pH, so it is not a buffer solution.
05

Option e: \(0.85\,\mathrm{M}\,\mathrm{H}_2\,\mathrm{NNH}_2 + 0.60\,\mathrm{M}\, \mathrm{H}_2\,\mathrm{NNH}_3\,\mathrm{NO}_3\)

Here, \(H_2NNH_2\) (hydrazine) is a weak base, and its conjugate acid is \(\mathrm{H}_{2} \mathrm{NNH}_{3}^{+}\) which is present in the form of its salt \(\mathrm{H}_{2}\mathrm{NNH}_{3}\mathrm{NO}_{3}\) in the solution. This system consists of a weak base and its conjugate acid in significant amounts, so it can be classified as a buffer solution. In conclusion, option c and option e can be classified as buffer solutions.

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Most popular questions from this chapter

Calculate the \(\mathrm{pH}\) of a solution prepared by mixing \(250 . \mathrm{mL}\) of \(0.174 m\) aqueous \(\mathrm{HF}\) (density \(=1.10 \mathrm{~g} / \mathrm{mL}\) ) with \(38.7 \mathrm{~g}\) of an aqueous solution that is \(1.50 \% \mathrm{NaOH}\) by mass (density \(=\) \(1.02 \mathrm{~g} / \mathrm{mL}) \cdot\left(K_{\mathrm{a}}\right.\) for \(\left.\mathrm{HF}=7.2 \times 10^{-4} .\right)\)

Calculate the \(\mathrm{pH}\) of each of the following buffered solutions. a. \(0.10 M\) acetic acid \(/ 0.25 M\) sodium acetate b. \(0.25 M\) acetic acid/0.10 \(M\) sodium acetate c. \(0.080 M\) acetic acid \(/ 0.20 M\) sodium acetate d. \(0.20 M\) acetic acid/0.080 \(M\) sodium acetate

Consider a buffered solution containing \(\mathrm{CH}_{3} \mathrm{NH}_{3} \mathrm{Cl}\) and \(\mathrm{CH}_{3} \mathrm{NH}_{2}\). Which of the following statements concerning this solution is(are) true? \(\left(K_{\mathrm{a}}\right.\) for \(\mathrm{CH}_{3} \mathrm{NH}_{3}^{+}=2.3 \times 10^{-11}\).) a. A solution consisting of \(0.10 \mathrm{M} \mathrm{CH}_{3} \mathrm{NH}_{3} \mathrm{Cl}\) and \(0.10 \mathrm{M}\) \(\mathrm{CH}_{3} \mathrm{NH}_{2}\) would have a higher buffering capacity than one containing \(1.0 \mathrm{M} \mathrm{CH}_{3} \mathrm{NH}_{3} \mathrm{Cl}\) and \(1.0 \mathrm{M} \mathrm{CH}_{3} \mathrm{NH}_{2}\). b. If \(\left[\mathrm{CH}_{3} \mathrm{NH}_{2}\right]>\left[\mathrm{CH}_{3} \mathrm{NH}_{3}{ }^{+}\right]\), then the \(\mathrm{pH}\) is larger than the \(\mathrm{p} K_{\mathrm{a}}\) value. c. Adding more \(\left[\mathrm{CH}_{3} \mathrm{NH}_{3} \mathrm{Cl}\right]\) to the initial buffer solution will decrease the \(\mathrm{pH}\). d. If \(\left[\mathrm{CH}_{3} \mathrm{NH}_{2}\right]<\left[\mathrm{CH}_{3} \mathrm{NH}_{3}{ }^{+}\right]\), then \(\mathrm{pH}<3.36\). e. If \(\left[\mathrm{CH}_{3} \mathrm{NH}_{2}\right]=\left[\mathrm{CH}_{3} \mathrm{NH}_{3}{ }^{+}\right]\), then \(\mathrm{pH}=10.64\).

A \(10.00-\mathrm{g}\) sample of the ionic compound NaA, where \(\mathrm{A}^{-}\) is the anion of a weak acid, was dissolved in enough water to make \(100.0 \mathrm{~mL}\) of solution and was then titrated with \(0.100 M\) HCl. After \(500.0 \mathrm{~mL}\) HCl was added, the \(\mathrm{pH}\) was \(5.00\). The experimenter found that \(1.00 \mathrm{~L}\) of \(0.100 \mathrm{M} \mathrm{HCl}\) was required to reach the stoichiometric point of the titration. a. What is the molar mass of \(\mathrm{NaA}\) ? b. Calculate the \(\mathrm{pH}\) of the solution at the stoichiometric point of the titration.

Derive an equation analogous to the Henderson-Hasselbalch equation but relating \(\mathrm{pOH}\) and \(\mathrm{p} K_{\mathrm{b}}\) of a buffered solution composed of a weak base and its conjugate acid, such as \(\mathrm{NH}_{3}\) and \(\mathrm{NH}_{4}^{+} .\)
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