/*! 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 9 The common ion effect for weak a... [FREE SOLUTION] | 91影视

91影视

Log out

Learning Materials

Features

Discover

Chapter 15: Problem 9

The common ion effect for weak acids is to significantly decrease the dissociation of the acid in water. Explain the common ion effect.

Short Answer

Expert verified
The common ion effect is the suppression of the dissociation of a weak electrolyte, like a weak acid, due to the presence of a common ion in the solution. A weak acid partially dissociates in water, reaching an equilibrium between undissociated acid and produced ions. For example, adding sodium acetate (CH鈧僀OONa) to a solution of acetic acid (CH鈧僀OOH) increases the concentration of the common ion CH鈧僀OO鈦�. According to Le Ch芒telier's principle, the system adjusts to reestablish equilibrium by shifting towards the formation of undissociated CH鈧僀OOH molecules, which decreases the dissociation of the weak acid.

Step by step solution

01

Define the Common Ion Effect

The common ion effect refers to the suppression of the dissociation of a weak electrolyte (like a weak acid or weak base) when another electrolyte containing a common ion is added to the solution. In other words, the presence of a common ion in the solution reduces the solubility of the weak electrolyte.
02

Describe Weak Acids

A weak acid is an acid that does not fully dissociate in water, meaning not all the molecules break apart into ions. Instead, a partial equilibrium is reached between the undissociated acid molecules and the produced ions. The level of dissociation in water for a given weak acid can be represented by an equilibrium constant, known as the acid dissociation constant (K_a). For example, considering a generic weak acid, HA: \[HA \rightleftharpoons H^+ + A^-\] The acid dissociation constant (K_a) is given by: \[K_a = \frac{[H^+][A^-]}{[HA]}\] Where [H^+], [A^-], and [HA] represent the equilibrium concentrations of the respective species.
03

Illustrate the Common Ion Effect with an Example

Consider the weak acid acetic acid, CH鈧僀OOH, which dissociates according to the following equation: \[CH_3COOH \rightleftharpoons H^+ + CH_3COO^-\] Now let's add a common ion, for example, sodium acetate (CH鈧僀OONa), which is a strong electrolyte and completely dissociates into its ions in water: \[CH_3COONa \rightarrow CH_3COO^- + Na^+\] As a result, the concentration of the common ion CH鈧僀OO鈦� increases. According to Le Ch芒telier's principle, the system will now adjust to counteract this change and reestablish the equilibrium. In this case, the reaction will shift towards the left, favoring the formation of undissociated CH鈧僀OOH molecules and decreasing the dissociation of the weak acid.
04

Explain the Role of Equilibrium and Le Ch芒telier's Principle

The common ion effect can be explained using the concepts of equilibrium and Le Ch芒telier's principle. In a weak acid dissociation reaction, an equilibrium exists between the undissociated acid and its dissociated ions. Le Ch芒telier's principle states that if an external stress, e.g., a change in concentration, is applied to a system at equilibrium, the system will adjust to minimize the effect of the stress and restore equilibrium. In the case of the common ion effect, the increased concentration of the common ion causes the equilibrium to shift towards the undissociated acid, thus reducing the dissociation.

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

Consider the titration of \(40.0 \mathrm{~mL}\) of \(0.200 \mathrm{M} \mathrm{HClO}_{4}\) by \(0.100\) \(M\) KOH. Calculate the \(\mathrm{pH}\) of the resulting solution after the following volumes of KOH have been added. a. \(0.0 \mathrm{~mL}\) d. \(80.0 \mathrm{~mL}\) b. \(10.0 \mathrm{~mL}\) e. \(100.0 \mathrm{~mL}\) c. \(40.0 \mathrm{~mL}\)

Calculate the \(\mathrm{pH}\) of a solution prepared by mixing \(250 . \mathrm{mL}\) of \(0.174 \mathrm{~m}\) 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 \(=1.02\) \(\mathrm{g} / \mathrm{mL}\) ). \(\left(K_{\mathrm{a}}\right.\) for \(\mathrm{HF}=7.2 \times 10^{-4}\).)

A few drops of each of the indicators shown in the accompanying table were placed in separate portions of a \(1.0 \mathrm{M}\) solution of a weak acid, \(\mathrm{HX}\). The results are shown in the last column of the table. What is the approximate \(\mathrm{pH}\) of the solution containing HX? Calculate the approximate value of \(K_{\alpha}\) for \(\mathrm{HX}\). \begin{tabular}{|lclcc|} \hline Indicator & Color of Hin & Color of \(\mathrm{In}^{-}\) & \(\mathrm{p} \boldsymbol{K}_{\mathrm{a}}\) of Hin & Color of \(1.0 \mathrm{M} \mathrm{HX}\) \\\ \hline Bromphenol blue & Yellow & Blue & \(4.0\) & Blue \\ Bromcresol purple & Yellow & Purple & \(6.0\) & Yellow \\ Bromcresol green & Yellow & Blue & \(4.8\) & Green \\ Alizarin & Yellow & Red & \(6.5\) & Yellow \\ \hline \end{tabular}

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.

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

Recommended explanations on Chemistry Textbooks

Physical Chemistry

Read Explanation

Chemical Reactions

Read Explanation

Chemical Analysis

Read Explanation

Organic Chemistry

Read Explanation

Making Measurements

Read Explanation

Ionic and Molecular Compounds

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.