/*! This file is auto-generated */ .wp-block-button__link{color:#fff;background-color:#32373c;border-radius:9999px;box-shadow:none;text-decoration:none;padding:calc(.667em + 2px) calc(1.333em + 2px);font-size:1.125em}.wp-block-file__button{background:#32373c;color:#fff;text-decoration:none} Problem 3 Mixing together solutions of ace... [FREE SOLUTION] | 91影视

91影视

Log out

Learning Materials

Features

Discover

Chapter 15: Problem 3

Mixing together solutions of acetic acid and sodium hydroxide can make a buffered solution. Explain. How does the amount of each solution added change the effectiveness of the buffer?

Short Answer

Expert verified
Mixing acetic acid (CH3COOH) and sodium hydroxide (NaOH) forms a buffered solution containing acetic acid and its conjugate base, acetate ions (CH3COO-). A buffer solution resists changes in pH when small amounts of acid or base are added. The effectiveness of a buffer solution depends on the concentrations of the weak acid and its conjugate base. The Henderson-Hasselbalch equation illustrates this relationship: \( pH = pK_a + log_{10} \frac{[CH_3COO^-]}{[CH_3COOH]} \). By adjusting the amount of acetic acid and sodium hydroxide added, we can control the pH of the resulting buffered solution. If the concentrations of acetic acid and acetate ions are nearly equal, the buffer solution will have a maximum buffering capacity, making it more effective at resisting changes in pH.

Step by step solution

01

Understand the concept of a buffer solution

A buffer solution is a solution that resists changes in pH when small amounts of an acid or a base are added to it. It consists of a mixture of a weak acid and its conjugate base, or a weak base and its conjugate acid.
02

Understand the reaction between acetic acid and sodium hydroxide

When acetic acid (CH3COOH), a weak acid, is mixed with sodium hydroxide (NaOH), a strong base, a neutralization reaction occurs, producing sodium acetate (CH3COONa) and water (H2O). The balanced equation for this reaction is: \( CH_3COOH + NaOH \rightarrow CH_3COONa + H_2O \) Sodium acetate, formed during this reaction, dissociates in water to form acetate ions (CH3COO-) and sodium ions (Na+). The acetate ion is the conjugate base of acetic acid. The acidic component of the buffer system is acetic acid, and the basic component is the acetate ion.
03

Formation of the buffer solution

When acetic acid and sodium hydroxide are mixed, the resulting solution contains both acetic acid and its conjugate base, acetate ions: \( CH_3COOH \leftrightarrow H^+ + CH_3COO^- \) Since both acidic and basic components are present in the solution, it forms a buffer system that resists changes in pH when small amounts of acid or base are added.
04

Effect of the amount of each solution on buffer effectiveness

The effectiveness of a buffer solution depends on the concentrations of the weak acid and its conjugate base. The Henderson-Hasselbalch equation describes the relationship between the pH of a buffer solution and the concentrations of the acidic and basic components: \( pH = pK_a + log_{10} \frac{[CH_3COO^-]}{[CH_3COOH]} \) According to this equation, the pH of the buffer solution depends on the ratio of the concentrations of acetate ions (CH3COO-) and acetic acid (CH3COOH). By adjusting the amount of acetic acid and sodium hydroxide added to the solution, we can control the ratio of these two components and, consequently, the pH of the resulting buffered solution. If the concentrations of acetic acid and acetate ions are nearly equal, the buffer solution will be most effective at resisting changes in pH, as it will have a maximum buffering capacity. If the concentration of one component is much higher than the other, the buffering capacity will be decreased, making the buffer solution less effective at resisting changes in pH.

Unlock Step-by-Step Solutions & Ace Your Exams!

  • Full Textbook Solutions

    Get detailed explanations and key concepts

  • Unlimited Al creation

    Al flashcards, explanations, exams and more...

  • Ads-free access

    To over 500 millions flashcards

  • Money-back guarantee

    We refund you if you fail your exam.

Over 30 million students worldwide already upgrade their learning with 91影视!

One App. One Place for Learning.

All the tools & learning materials you need for study success - in one app.

Get started for free

Most popular questions from this chapter

Calculate the mass of sodium acetate that must be added to \(500.0 \mathrm{~mL}\) of \(0.200 \mathrm{M}\) acetic acid to form a \(\mathrm{pH}=5.00\) buffer solution.

Methyl red has the following structure: CN(C)c1ccc(N=Nc2ccccc2C(=O)O)cc1 It undergoes a color change from red to yellow as a solution gets more basic. Calculate an approximate \(\mathrm{pH}\) range for which methyl red is useful. What is the color change and the \(\mathrm{pH}\) at the color change when a weak acid is titrated with a strong base using methyl red as an indicator? What is the color change and the \(\mathrm{pH}\) at the color change when a weak base is titrated with a strong acid using methyl red as an indicator? For which of these two types of titrations is methyl red a possible indicator?

Lactic acid is a common by-product of cellular respiration and is often said to cause the "burn" associated with strenuous activity. A \(25.0-\mathrm{mL}\) sample of \(0.100 M\) lactic acid \(\left(\mathrm{HC}_{3} \mathrm{H}_{5} \mathrm{O}_{3}, \mathrm{p} K_{\mathrm{a}}=\right.\) \(3.86\) ) is titrated with \(0.100 M\) NaOH solution. Calculate the \(\mathrm{pH}\) after the addition of \(0.0 \mathrm{~mL}, 4.0 \mathrm{~mL}, 8.0 \mathrm{~mL}, 12.5 \mathrm{~mL}, 20.0 \mathrm{~mL}\), \(24.0 \mathrm{~mL}, 24.5 \mathrm{~mL}, 24.9 \mathrm{~mL}, 25.0 \mathrm{~mL}, 25.1 \mathrm{~mL}, 26.0 \mathrm{~mL}, 28.0 \mathrm{~mL}\) and \(30.0 \mathrm{~mL}\) of the \(\mathrm{NaOH}\). Plot the results of your calculations as pH versus milliliters of \(\mathrm{NaOH}\) added.

A \(10.00-g\) sample of the ionic compound \(\mathrm{NaA}\), 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 \(0.100 \mathrm{M} \mathrm{HCl}\). 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 NaA? b. Calculate the \(\mathrm{pH}\) of the solution at the stoichiometric point of the titration.

A \(0.400 M\) solution of ammonia was titrated with hydrochloric acid to the equivalence point, where the total volume was \(1.50\) times the original volume. At what \(\mathrm{pH}\) does the equivalence point occur?
See all solutions

Recommended explanations on Chemistry Textbooks

Ionic and Molecular Compounds

Read Explanation

Chemistry Branches

Read Explanation

Physical Chemistry

Read Explanation

Making Measurements

Read Explanation

Chemical Analysis

Read Explanation

Nuclear Chemistry

Read Explanation
View all explanations

What do you think about this solution?

We value your feedback to improve our textbook solutions.

Study anywhere. Anytime. Across all devices.