Textbook Question
Predict how each molecule or ion would act, in the Brønsted-Lowry sense, in aqueous solution by writing 'acid,' 'base,' 'both,' or 'neither' on the line provided. (b) Prozac
Ch.16 - Acid-Base Equilibria
Chapter 16, Problem 106
Oxalic acid H2C2O4 is a diprotic acid. By using data in Appendix D as needed, determine whether each of the following statements is true: (a) H2C2O4 can serve as both a Brønsted–Lowry acid and a Brønsted–Lowry base. (b) C2O4²⁻ is the conjugate base of HC2O4⁻. (c) An aqueous solution of the strong electrolyte KHC2O4 will have a pH of 6 or 7.
Verified step by step guidance1
Step 1: Understand the nature of oxalic acid (H2C2O4) as a diprotic acid, which means it can donate two protons (H⁺ ions) in solution. Write the dissociation reactions: H2C2O4 ⇌ H⁺ + HC2O4⁻ and HC2O4⁻ ⇌ H⁺ + C2O4²⁻.
Step 2: For statement (a), consider the definitions of Brønsted–Lowry acids and bases. A Brønsted–Lowry acid donates a proton, while a Brønsted–Lowry base accepts a proton. Analyze if H2C2O4 can act as both by looking at its ability to donate and accept protons in the dissociation reactions.
Step 3: For statement (b), identify the conjugate base of HC2O4⁻. Recall that the conjugate base is formed when an acid donates a proton. From the dissociation reaction HC2O4⁻ ⇌ H⁺ + C2O4²⁻, determine if C2O4²⁻ is indeed the conjugate base of HC2O4⁻.
Step 4: For statement (c), consider the nature of KHC2O4 as a salt formed from a strong base (KOH) and a weak acid (HC2O4⁻). Analyze the hydrolysis of HC2O4⁻ in water to determine if the solution will be acidic, neutral, or basic, and thus if the pH will be around 6 or 7.
Step 5: Use Appendix D to find relevant pKa values for the dissociation of H2C2O4 and compare them to typical pH values to support your analysis of each statement. This will help confirm the acidity or basicity of the solutions involved.
Key Concepts
Here are the essential concepts you must grasp in order to answer the question correctly.
Brønsted–Lowry Theory
The Brønsted–Lowry theory defines acids as proton donors and bases as proton acceptors. This concept is essential for understanding the behavior of substances in acid-base reactions. For example, oxalic acid (H2C2O4) can donate protons to become HC2O4⁻, thus acting as an acid, while HC2O4⁻ can accept a proton to revert to H2C2O4, demonstrating its ability to function as a base.
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02:08
Bronsted-Lowry Acid-Base Theory
Conjugate Acid-Base Pairs
Conjugate acid-base pairs consist of two species that differ by the presence of a proton (H⁺). In this context, HC2O4⁻ is the conjugate acid of C2O4²⁻, as it can donate a proton to form C2O4²⁻. Understanding these pairs is crucial for analyzing the relationships between acids and bases in chemical reactions.
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01:30Conjugate Acid-Base Pairs
pH and Strong Electrolytes
pH is a measure of the acidity or basicity of a solution, with lower values indicating higher acidity. Strong electrolytes, like KHC2O4, dissociate completely in solution, affecting the pH. Since KHC2O4 is a salt derived from a weak acid (HC2O4⁻) and a strong base (KOH), its solution typically has a pH around neutral (6 or 7), which is important for predicting the behavior of the solution.
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02:50Electrolytes and Strong Acids
Related Practice
Textbook Question
Calculate the pH of a solution made by adding 2.50 g of lithium oxide 1Li2O2 to enough water to make 1.500 L of solution.
Textbook Question
Butyric acid is responsible for the foul smell of rancid butter. The pKa of butyric acid is 4.84. (a) Calculate the pKb for the butyrate ion.
Textbook Question
Butyric acid is responsible for the foul smell of rancid butter. The pKa of butyric acid is 4.84. (c) Calculate the pH of a 0.050 M solution of sodium butyrate.