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Chapter 17: Problem 98

Write equations to show how each of the following buffer solutions reacts with a small added amount of a strong acid or a strong base: (a) HCOOH-KHCOO; (b) \(\mathrm{C}_{6} \mathrm{H}_{5} \mathrm{NH}_{2}-\mathrm{C}_{6} \mathrm{H}_{5} \mathrm{NH}_{3}^{+} \mathrm{Cl}^{-}\) (c) \(\mathrm{KH}_{2} \mathrm{PO}_{4}-\mathrm{Na}_{2} \mathrm{HPO}_{4}\)

Short Answer

Expert verified
In the presence of a strong acid, the base component of the buffer reacts to form its conjugate. When a strong base is present, the acid component reacts to form water and its conjugate.

Step by step solution

01

Reaction with strong acid

When an acid (represented here as H+) is added to a buffer solution, the base component of the buffer reacts to form its conjugate. Therefore, the reactions are: (a) \( KHCOO^- + H^{+} \rightarrow HCOOH \)(b) \( C_6H_5NH_2 + H^{+} \rightarrow C_6H_5NH_3^+ \)(c) \( Na_2HPO_4^- + H^{+} \rightarrow KH_2PO_4 \)
02

Reaction with strong base

When a base (represented here as OH-) is added to the solution, the acid component of the buffer reacts to form water and its conjugate. Therefore, the reactions are: (a) \( HCOOH + OH^- \rightarrow H_2O + HCOO^- \)(b) \( C_6H_5NH_3^+Cl^- + OH^- \rightarrow H_2O + C_6H_5NH_2 \)(c) \( KH_2PO_4 + OH^- \rightarrow H_2O + Na_2HPO_4^- \)

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Most popular questions from this chapter

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?

Calculate the pH of the buffer formed by mixing equal volumes \(\left[\mathrm{C}_{2} \mathrm{H}_{5} \mathrm{NH}_{2}\right]=1.49 \mathrm{M} \quad\) with \(\quad\left[\mathrm{HClO}_{4}\right]=\) 1.001 M. \(K_{\mathrm{b}}=4.3 \times 10^{-4}\)

Is a solution that is \(0.10 \mathrm{M} \mathrm{Na}_{2} \mathrm{S}(\mathrm{aq})\) likely to be acidic, basic, or pH neutral? Explain.

Two solutions are mixed: \(100.0 \mathrm{mL}\) of \(\mathrm{HCl}(\mathrm{aq})\) with \(\mathrm{pH} 2.50\) and \(100.0 \mathrm{mL}\) of \(\mathrm{NaOH}(\text { aq) with } \mathrm{pH} 11.00\) What is the pH of the resulting solution?

In the use of acid-base indicators, (a) Why is it generally sufficient to use a single indicator in an acid-base titration, but often necessary to use several indicators to establish the approximate pH of a solution? (b) Why must the quantity of indicator used in a titration be kept as small as possible?
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