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Chapter 14: Problem 18

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

Short Answer

Expert verified
Mixtures c. (\(0.50 M\) HOCl \(+0.35 M\) KOCl) and e. (\(0.85 M \mathrm{H}_{2} \mathrm{NNH}_{2}+0.60 \mathrm{M} \mathrm{H}_{2} \mathrm{NNH}_{3} \mathrm{NO}_{3}\)) can be classified as buffer solutions.

Step by step solution

01

- Identifying Mixture Components

We will look at each of the mixtures provided and identify whether they contain a weak acid, a weak base, and their conjugate pairs. a. \(0.25 M\) HBr \(+0.25 M\) HOBr b. \(0.15 M \mathrm{HClO}_{4}+0.20 \mathrm{M} \mathrm{RbOH}\) c. \(0.50 M\) HOCl \(+0.35 M\) KOCl d. \(0.70 M \mathrm{KOH}+0.70 \mathrm{M} \mathrm{HONH}_{2}\) e. \(0.85 M \mathrm{H}_{2} \mathrm{NNH}_{2}+0.60 \mathrm{M} \mathrm{H}_{2} \mathrm{NNH}_{3} \mathrm{NO}_{3}\)
02

- Evaluating Mixture Components

We will evaluate each of the mixtures based on whether they contain a weak acid/base and their conjugate pairs. a. HBr is a strong acid, and HOBr is a weak acid. However, there isn't a conjugate base present, so this isn't a buffer solution. b. \(\mathrm{HClO}_{4}\) is a strong acid, and \(\mathrm{RbOH}\) is a strong base. Neither of them is a weak acid/base with a conjugate pair. Thus, this isn't a buffer solution. c. HOCl is a weak acid, and its conjugate base is \(\mathrm{OCl^{-}}\) which is present as KOCl, making this a buffer solution. d. \(\mathrm{KOH}\) is a strong base, and \(\mathrm{HONH}_2\) is a weak base. However, there isn't a conjugate acid present, so this isn't a buffer solution. e. \(\mathrm{H}_{2}\mathrm{NNH}_{2}\) is a weak base and its conjugate acid is \(\mathrm{H}_{2}\mathrm{NNH}_{3}^{+}\), which is present in the form of \(\mathrm{H}_{2}\mathrm{NNH}_{3}\mathrm{NO}_{3}\). This mixture is a buffer solution.
03

- Conclusion

After evaluating each of the mixtures, we found that only mixtures c. (\(0.50 M\) HOCl \(+0.35 M\) KOCl) and e. (\(0.85 M \mathrm{H}_{2} \mathrm{NNH}_{2}+0.60 \mathrm{M} \mathrm{H}_{2} \mathrm{NNH}_{3} \mathrm{NO}_{3}\)) can be classified as buffer solutions.

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Key Concepts

These are the key concepts you need to understand to accurately answer the question.

Acid-Base Equilibrium
The concept of acid-base equilibrium is central to understanding buffer solutions. It involves the balance that exists between acids and bases in a solution. This delicate equilibrium can be represented by the equation for the dissociation of a weak acid or base:

For a weak acid, HA, the equilibrium with its dissociated ions is represented as: \[ HA \rightleftharpoons H^+ + A^- \]
Similarly, for a weak base, BOH, it dissociates into: \[ BOH \rightleftharpoons B^+ + OH^- \]
In both cases, the equilibrium lies more towards the left, which means these substances do not completely dissociate in solution.
An effective buffer solution contains a weak acid or base along with its conjugate pair at significant concentrations, maintaining this equilibrium even when small amounts of strong acids or bases are added.
Conjugate Acid-Base Pairs
When we speak of conjugate acid-base pairs, we refer to a pair of substances that differ by the presence of a single hydrogen ion. A conjugate acid can donate a hydrogen ion, whereas its conjugate base can accept one. This pair is found in a buffer solution where the acid/base and its conjugate ensure the pH remains relatively stable.

For instance, acetic acid (\( CH_3COOH \)) and its conjugate base, acetate (\( CH_3COO^- \)), make a common buffer pair. When acetic acid loses a hydrogen ion, it becomes acetate, and vice versa. This transformation between the two under changing pH conditions is what keeps the solution buffered. Understanding the relation between these pairs is crucial for grasping how buffers work and predicting the outcomes of acid-base reactions.
Weak Acids and Bases
The presence of weak acids or bases is a defining characteristic of buffer solutions. These substances do not completely ionize in solution, making them ideal for creating a buffer. This incomplete ionization allows them to resist changes in pH because they can react with any added strong acid or base to maintain the equilibrium.

In a buffer solution, a weak acid will have a corresponding conjugate base, usually in the form of a salt, that helps to

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Most popular questions from this chapter

Which of the following mixtures would result in a buffered solution when 1.0 L of each of the two solutions are mixed? a. \(0.2 M\) HNO and \(0.4 M \mathrm{NaNO}_{3}\) b. \(0.2 M \mathrm{HNO}_{3}\) and \(0.4 \mathrm{M} \mathrm{HF}\) c. \(0.2 \mathrm{M} \mathrm{HNO}_{3}\) and \(0.4 \mathrm{M} \mathrm{NaF}\) d. \(0.2 M\) HNO \(_{3}\) and \(0.4 M\) NaOH

Consider the titration of \(100.0 \mathrm{mL}\) of \(0.100 \mathrm{M}\) HCN by \(0.100 M \mathrm{KOH}\) at \(25^{\circ} \mathrm{C} .\left(K_{\mathrm{a}} \text { for } \mathrm{HCN}=6.2 \times 10^{-10} .\right)\) a. Calculate the \(\mathrm{pH}\) after \(0.0 \mathrm{mL}\) of \(\mathrm{KOH}\) has been added. b. Calculate the \(\mathrm{pH}\) after \(50.0 \mathrm{mL}\) of \(\mathrm{KOH}\) has been added. c. Calculate the \(\mathrm{pH}\) after \(75.0 \mathrm{mL}\) of \(\mathrm{KOH}\) has been added. d. Calculate the \(\mathrm{pH}\) at the equivalence point. e. Calculate the pH after 125 mL of KOH has been added.

Calculate the \(\mathrm{pH}\) at the halfway point and at the equivalence point for each of the following titrations. a. \(100.0 \mathrm{mL}\) of \(0.10 \mathrm{M} \mathrm{HC}_{7} \mathrm{H}_{5} \mathrm{O}_{2}\left(K_{\mathrm{a}}=6.4 \times 10^{-5}\right)\) titrated by 0.10 \(M \mathrm{NaOH}\) b. \(100.0 \mathrm{mL}\) of \(0.10 \mathrm{M} \mathrm{C}_{2} \mathrm{H}_{5} \mathrm{NH}_{2}\left(K_{\mathrm{b}}=5.6 \times 10^{-4}\right)\) titrated by 0.20 \(M \mathrm{HNO}_{3}\) c. \(100.0 \mathrm{mL}\) of \(0.50 \mathrm{M}\) HCl titrated by \(0.25 \mathrm{M} \mathrm{NaOH}\)

Sketch a pH curve for the titration of a weak acid (HA) with a strong base (NaOH). List the major species, and explain how you would go about calculating the pH of the solution at various points, including the halfway point and the equivalence point.

Amino acids are the building blocks for all proteins in our bodies. A structure for the amino acid alanine is All amino acids have at least two functional groups with acidic or basic properties. In alanine, the carboxylic acid group has \(K_{\mathrm{a}}=4.5 \times 10^{-3}\) and the amino group has \(K_{\mathrm{b}}=\) \(7.4 \times 10^{-5} .\) Because of the two groups with acidic or basic properties, three different charged ions of alanine are possible when alanine is dissolved in water. Which of these ions would predominate in a solution with \(\left[\mathrm{H}^{+}\right]=1.0\) \(\mathrm{M} ?\) In a solution with \(\left[\mathrm{OH}^{-}\right]=1.0\) \(\mathrm {M} ?\)
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