Textbook Question
Calculate the percent ionization of propionic acid (C2H5COOH) in solutions of each of the following concentrations (Ka is given in Appendix D): (a) 0.250 M (b) 0.0800 M (c) 0.0200 M
Ch.16 - Acid-Base Equilibria
Chapter 16, Problem 66
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?
Verified step by step guidance1
Identify the relevant acid-dissociation constants (Ka1 and Ka2) for tartaric acid from Table 16.3, as tartaric acid is a diprotic acid.
Write the balanced chemical equations for the stepwise dissociation of tartaric acid: \( \text{H}_2\text{C}_4\text{H}_4\text{O}_6 \rightleftharpoons \text{H}^+ + \text{HC}_4\text{H}_4\text{O}_6^- \) and \( \text{HC}_4\text{H}_4\text{O}_6^- \rightleftharpoons \text{H}^+ + \text{C}_4\text{H}_4\text{O}_6^{2-} \).
Set up the equilibrium expressions for each dissociation step using the acid-dissociation constants: \( K_{a1} = \frac{[\text{H}^+][\text{HC}_4\text{H}_4\text{O}_6^-]}{[\text{H}_2\text{C}_4\text{H}_4\text{O}_6]} \) and \( K_{a2} = \frac{[\text{H}^+][\text{C}_4\text{H}_4\text{O}_6^{2-}]}{[\text{HC}_4\text{H}_4\text{O}_6^-]} \).
Assume that the first dissociation is the primary contributor to the \([\text{H}^+]\) concentration, and use the initial concentration of tartaric acid (0.250 M) to solve for \([\text{H}^+]\) using \( K_{a1} \).
Use the \([\text{H}^+]\) concentration from the first dissociation to calculate the \([\text{C}_4\text{H}_4\text{O}_6^{2-}]\) concentration from the second dissociation using \( K_{a2} \), and then calculate the pH as \( \text{pH} = -\log[\text{H}^+] \).
Key Concepts
Here are the essential concepts you must grasp in order to answer the question correctly.
Acid-Base Equilibria
Acid-base equilibria involve the dissociation of acids in solution, which can be described by their acid-dissociation constants (Ka). For tartaric acid, a diprotic acid, two dissociation steps must be considered, each with its own Ka value. Understanding how these equilibria shift in response to concentration changes is crucial for calculating pH and ion concentrations.
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Arrhenius Acids and Bases
pH Calculation
The pH of a solution is a measure of its acidity, defined as the negative logarithm of the hydrogen ion concentration (pH = -log[H⁺]). In the case of tartaric acid, the pH can be calculated using the concentrations of the dissociated ions and the equilibrium constants. This requires applying the principles of equilibrium and stoichiometry to find the concentration of H⁺ ions in the solution.
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Assumptions in Calculations
When performing calculations involving weak acids like tartaric acid, certain assumptions may simplify the process, such as neglecting the contribution of H⁺ from water or assuming that the change in concentration due to dissociation is negligible. These approximations can affect the accuracy of the results, so it is important to evaluate their validity based on the initial concentrations and the strength of the acid.
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Related Practice
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?
Textbook Question
Consider the base hydroxylamine, NH2OH. (a) What is the conjugate acid of hydroxylamine?
Textbook Question
The hypochlorite ion, ClO-, acts as a weak base. (a) Is ClO- a stronger or weaker base than hydroxylamine?