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Chapter 17: Problem 14

What is the pH of a solution prepared by dissolving \(8.50 \mathrm{g}\) of aniline hydrochloride \(\left(\mathrm{C}_{6} \mathrm{H}_{5} \mathrm{NH}_{3}^{+} \mathrm{Cl}^{-}\right)\) in \(750 \mathrm{mL}\) of \(0.215 \mathrm{M}\) aniline, \(\left(\mathrm{C}_{6} \mathrm{H}_{5} \mathrm{NH}_{2}\right) ?\) Would this solution be an effective buffer? Explain.

Short Answer

Expert verified
First calculate the moles of aniline and aniline hydrochloride and their respective concentrations. Then, use these values in the Henderson-Hasselbalch equation along with the pKa of aniline to find the pH of the solution. Lastly, use the buffer capacity rule to determine whether or not the solution could act as a good buffer.

Step by step solution

01

Calculate the moles of aniline hydrochloride

Firstly, calculate the number of moles of aniline hydrochloride dissolved. Use the given mass and the molar mass of aniline hydrochloride. According to the periodic table, the molar mass of aniline hydrochloride (represented by C6H5NH3+Cl-) is approximately 129.6 g/mol. Thus, the moles of aniline hydrochloride \( n_{(aniline \, hydrochloride)} = \frac{8.50 \, g}{129.6 \, g/mol}\)
02

Calculate the moles of aniline

Next, calculate how many moles of aniline are in the solution prior to the addition of aniline hydrochloride. This is done by multiplying the concentration of aniline given by the volume in litres: \( n_{(aniline)} = M_{(aniline)} * V_{(aniline)} = 0.215 \, M * 0.75 \, L \)
03

Determine acid and base concentrations

Then, calculate the concentrations of the aniline hydrochloride and aniline. The total volume should be used to calculate these concentrations, which will be the volume of the aniline solution plus the volume of the aniline hydrochloride solution. However, since no volume is given for the aniline hydrochloride solution, assume that its volume is negligible and calculate with only the volume of the aniline solution. \( [base] = \frac{n_{(aniline)}}{V} \) , \( [acid] = \frac{n_{(aniline \, hydrochloride)}}{V} \)
04

Use the Henderson-Hasselbalch equation

Now, input the values obtained in the previous steps along with the pKa value of aniline into the Henderson-Hasselbach equation to calculate the pH of the buffer solution. The pKa of aniline is 4.63, and the Henderson-Hasselbalch equation is: \( pH = pKa + log \left(\frac{[base]}{[acid]}\right) \)
05

Evaluate buffer capacity

Finally, evaluate whether the solution could act as an effective buffer by using the buffer capacity rule. This rule states that a solution can act as a good buffer if the ratio \([base]/[acid]\) falls between 0.1 and 10. Calculate this ratio and compare it to the given values.

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Most popular questions from this chapter

To convert \(\mathrm{NH}_{4}^{+}(\text {aq })\) to \(\mathrm{NH}_{3}(\mathrm{aq}),\) (a) add \(\mathrm{H}_{3} \mathrm{O}^{+}\) (b) raise the \(\mathrm{pH} ;\) (c) add \(\mathrm{KNO}_{3}(\mathrm{aq}) ;\) (d) add \(\mathrm{NaCl}\).

What stoichiometric concentration of the indicated substance is required to obtain an aqueous solution with the pH value shown: (a) aniline, \(\mathrm{C}_{6} \mathrm{H}_{5} \mathrm{NH}_{2}\), for \(\mathrm{pH}=8.95 ;(\mathrm{b}) \mathrm{NH}_{4} \mathrm{Cl}\) for \(\mathrm{pH}=5.12 ?\)

Calculate the \(\mathrm{pH}\) at the points in the titration of \(25.00 \mathrm{mL}\) of \(0.132 \mathrm{M} \mathrm{HNO}_{2}\) when (a) \(10.00 \mathrm{mL}\) and (b) 20.00 mL of 0.116 M \(\mathrm{HNO}_{2}\) when \((\mathrm{a}) 10.00 \mathrm{mL}\) and (b) \(20.00 \mathrm{mL}\) of \(0.116 \mathrm{M}\) NaOH have been added. For \(\mathrm{HNO}_{2}, K_{\mathrm{a}}=7.2 \times 10^{-4}\) \(\mathrm{HNO}_{2}+\mathrm{OH}^{-} \longrightarrow \mathrm{H}_{2} \mathrm{O}+\mathrm{NO}_{2}^{-}\)

Consider a solution containing two weak monoprotic acids with dissociation constants \(K_{\mathrm{HA}}\) and \(K_{\mathrm{HB}}\). Find the charge balance equation for this system, and use it to derive an expression that gives the concentration of \(\mathrm{H}_{3} \mathrm{O}^{+}\) as a function of the concentrations of \(\mathrm{HA}\) and HB and the various constants.

You are asked to prepare a buffer solution with a pH of 3.50. The following solutions, all \(0.100 \mathrm{M},\) are available to you: HCOOH, CH \(_{3} \mathrm{COOH}, \mathrm{H}_{3} \mathrm{PO}_{4}, \mathrm{NaCHOO}\) \(\mathrm{NaCH}_{3} \mathrm{COO},\) and \(\mathrm{NaH}_{2} \mathrm{PO}_{4} . \quad\) Describe how you would prepare this buffer solution. [Hint: What volumes of which solutions would you use?
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