Textbook Question
The AIDS drug zalcitabine (also known as ddC) is a weak base with a pKb of 9.8. What percentage of the base is protonated in an aqueous zalcitabine solution containing 565 mg/L?
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Ch.17 - Acids and Bases
Tro6th EditionChemistry: A Molecular ApproachISBN: 9780137832217
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Table of contents
- Ch.1 - Matter, Measurement & Problem Solving90
- Ch.2 - Atoms & Elements94
- Ch.3 - Molecules and Compounds103
- Ch.4 - Chemical Reactions and Chemical Quantities46
- Ch.5 - Introduction to Solutions and Aqueous Solutions65
- Ch.6 - Gases100
- Ch.7 - Thermochemistry85
- Ch.8 - The Quantum-Mechanical Model of the Atom51
- Ch.9 - Periodic Properties of the Elements79
- Ch.10 - Chemical Bonding I: The Lewis Model70
- Ch.11 - Chemical Bonding II: Molecular Shapes, VSEPR & MO Theory68
- Ch.12 - Liquids, Solids & Intermolecular Forces44
- Ch.13 - Solids & Modern Materials42
- Ch.14 - Solutions77
- Ch.15 - Chemical Kinetics88
- Ch.16 - Chemical Equilibrium57
- Ch.17 - Acids and Bases121
- Ch.18 - Aqueous Ionic Equilibrium126
- Ch.19 - Free Energy & Thermodynamics80
- Ch.20 - Electrochemistry87
- Ch.21 - Radioactivity & Nuclear Chemistry61
- Ch.22 - Organic Chemistry114
Chapter 17, Problem 139
Common aspirin is acetylsalicylic acid, which has the structure shown here and a pKa of 3.5. Calculate the pH of a solution in which one normal adult dose of aspirin (6.5 * 10^2 mg) is dissolved in 8.0 ounces of water.
Verified step by step guidance1
Convert the mass of aspirin from milligrams to grams by dividing by 1000.
Calculate the number of moles of aspirin using its molar mass (180.16 g/mol for acetylsalicylic acid).
Convert the volume of water from ounces to liters (1 ounce = 0.0295735 liters).
Determine the concentration of the aspirin solution in moles per liter (Molarity) by dividing the number of moles by the volume in liters.
Use the Henderson-Hasselbalch equation, \( \text{pH} = \text{pKa} + \log \left( \frac{[A^-]}{[HA]} \right) \), assuming that the concentration of the ionized form \([A^-]\) is equal to the concentration of the aspirin solution, to calculate the pH.
Key Concepts
Here are the essential concepts you must grasp in order to answer the question correctly.
pKa and pH Relationship
The pKa value of an acid indicates the pH at which half of the acid is dissociated into its conjugate base. For acetylsalicylic acid, with a pKa of 3.5, this means that at pH 3.5, the concentrations of the acid and its conjugate base are equal. Understanding this relationship is crucial for calculating the pH of a solution containing the acid.
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02:09
pH and pOH Calculations
Normality and Molarity
Normality (N) is a measure of concentration equivalent to molarity (M) but accounts for the reactive capacity of a solute. In the case of aspirin, knowing its molecular weight allows us to convert the mass of the dose into moles, which can then be used to determine the molarity of the solution. This is essential for calculating the resulting pH.
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01:13Molarity
Dissociation of Weak Acids
Weak acids, like acetylsalicylic acid, do not fully dissociate in solution. Instead, they establish an equilibrium between the undissociated acid and its ions. The extent of this dissociation can be described using the acid dissociation constant (Ka), which is derived from the pKa. This concept is vital for determining the concentration of hydrogen ions in the solution, which directly affects the pH.