Skip to main content
Ch.16 - Acid-Base Equilibria
Chapter 16, Problem 64a

Calculate the percent ionization of propionic acid 1C2H5COOH2 in solutions of each of the following concentrations 1Ka is given in Appendix D): (a) 0.250 M

Verified step by step guidance
1
Identify the chemical equation for the ionization of propionic acid: \( \text{C}_2\text{H}_5\text{COOH} \rightleftharpoons \text{C}_2\text{H}_5\text{COO}^- + \text{H}^+ \).
Write the expression for the acid dissociation constant \( K_a \) using the concentrations of the products and reactants: \( K_a = \frac{[\text{C}_2\text{H}_5\text{COO}^-][\text{H}^+]}{[\text{C}_2\text{H}_5\text{COOH}]} \).
Set up an ICE (Initial, Change, Equilibrium) table to determine the equilibrium concentrations of the species involved.
Assume that the initial concentration of \( \text{C}_2\text{H}_5\text{COOH} \) is 0.250 M and that the change in concentration due to ionization is \( x \).
Solve for \( x \) using the \( K_a \) expression and calculate the percent ionization using the formula: \( \text{Percent Ionization} = \left( \frac{x}{0.250} \right) \times 100 \).

Verified Solution

Video duration:
3m
This video solution was recommended by our tutors as helpful for the problem above.
Was this helpful?

Key Concepts

Here are the essential concepts you must grasp in order to answer the question correctly.

Ionization of Weak Acids

Weak acids, like propionic acid, do not fully dissociate in solution. Instead, they establish an equilibrium between the undissociated acid and its ions. The extent of this ionization is characterized by the acid dissociation constant (Ka), which quantifies the strength of the acid and its ability to donate protons (H+) in solution.
Recommended video:
Guided course
01:15
Calculating Percent Ionization of Weak Acids

Percent Ionization

Percent ionization is a measure of the degree to which an acid dissociates in solution, expressed as a percentage. It is calculated using the formula: (concentration of ionized acid / initial concentration of acid) × 100%. This value helps to understand the strength of the acid in a given concentration and how it changes with dilution.
Recommended video:
Guided course
06:08
Percent Ionization Example

Equilibrium Calculations

To calculate percent ionization, one must set up an equilibrium expression based on the initial concentration of the acid and the change in concentration due to ionization. This involves using the Ka value to find the concentration of ions at equilibrium, allowing for the determination of how much of the acid has ionized relative to its initial concentration.
Recommended video:
Guided course
02:38
Equilibrium Constant Calculation
Related Practice
Open Question
Saccharin, a sugar substitute, is a weak acid with pKa = 2.32 at 25 °C. It ionizes in aqueous solution as follows: HNC7H4SO31(aq) ⇌ H+(aq) + NC7H4SO3-(aq). What is the pH of a 0.10 M solution of this substance?
Open Question
The active ingredient in aspirin is acetylsalicylic acid 1HC9H7O42, a monoprotic acid with Ka = 3.3 * 10^-4 at 25 °C. What is the pH of a solution obtained by dissolving two extra-strength aspirin tablets, each containing 500 mg of acetylsalicylic acid, in 250 mL of water?
Open Question
Calculate the percent ionization of hydrazoic acid (HN3) in solutions of each of the following concentrations (Ka is given in Appendix D): (a) 0.400 M, (b) 0.100 M, (c) 0.0400 M.
Textbook Question

Citric acid, which is present in citrus fruits, is a triprotic acid (Table 16.3). (a) Calculate the pH of a 0.040 M solution of citric acid. (b) Did you have to make any approximations or assumptions in completing your calculations? (c) Is the concentration of citrate ion 1C6H5O7 3-2 equal to, less than, or greater than the H+ ion concentration?

1670
views
Open Question
Tartaric acid is found in many fruits, including grapes, and is partially responsible for the dry texture of certain wines. Calculate the pH and the tartrate ion C4H4O6²⁻ concentration for a 0.250 M solution of tartaric acid, for which the acid-dissociation constants are listed in Table 16.3. Did you have to make any approximations or assumptions in your calculation?
Textbook Question

Consider the base hydroxylamine, NH2OH. (a) What is the conjugate acid of hydroxylamine?

1199
views