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Ch.17 - Aqueous Ionic Equilibrium
Tro - Chemistry: A Molecular Approach 4th Edition
Tro4th EditionChemistry: A Molecular ApproachISBN: 9780134112831Not the one you use?Change textbook
All textbooksTro 4th EditionCh.17 - Aqueous Ionic EquilibriumProblem 53d
Chapter 17, Problem 53d

Determine whether or not the mixing of each pair of solutions results in a buffer. d. 175.0 mL of 0.10 M NH3; 150.0 mL of 0.12 M NaOH

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1
Identify the components needed for a buffer solution: a weak acid and its conjugate base, or a weak base and its conjugate acid.
Recognize that NH3 (ammonia) is a weak base and NaOH (sodium hydroxide) is a strong base.
Understand that a buffer solution requires a weak base and its conjugate acid, not a strong base.
Calculate the moles of NH3: \( \text{moles of NH3} = 0.175 \text{ L} \times 0.10 \text{ M} \).
Calculate the moles of NaOH: \( \text{moles of NaOH} = 0.150 \text{ L} \times 0.12 \text{ M} \).

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Key Concepts

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

Buffer Solutions

A buffer solution is a system that resists changes in pH upon the addition of small amounts of acid or base. It typically consists of a weak acid and its conjugate base or a weak base and its conjugate acid. Buffers are crucial in maintaining stable pH levels in various chemical and biological processes.
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Weak Bases and Strong Bases

Weak bases, like ammonia (NH3), partially ionize in solution, establishing an equilibrium between the base and its conjugate acid (NH4+). In contrast, strong bases, such as sodium hydroxide (NaOH), fully dissociate in solution, leading to a significant increase in hydroxide ions (OH-). Understanding the behavior of these bases is essential for determining buffer capacity.
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Concentration and Volume in Solution Chemistry

The concentration of a solution, expressed in molarity (M), indicates the amount of solute per liter of solution. When mixing solutions, the total volume and the concentrations of the individual components must be considered to determine the final concentrations of the resulting mixture. This is important for assessing whether the mixture can act as a buffer.
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Related Practice
Textbook Question

Determine whether or not the mixing of each pair of solutions results in a buffer. a. 75.0 mL of 0.10 M HF; 55.0 mL of 0.15 M NaF b. 150.0 mL of 0.10 M HF; 135.0 mL of 0.175 M HCl d. 125.0 mL of 0.15 M CH3NH2; 120.0 mL of 0.25 M CH3NH3Cl

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Textbook Question

Determine whether or not the mixing of each pair of solutions results in a buffer. b. 50.0 mL of 0.10 M HCl; 35.0 mL of 0.150 M NaOH

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Textbook Question

Determine whether or not the mixing of each pair of solutions results in a buffer. c. 165.0 mL of 0.10 M HF; 135.0 mL of 0.050 M KOH

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Textbook Question

Determine whether or not the mixing of each pair of solutions results in a buffer. a. 100.0 mL of 0.10 M NH3; 100.0 mL of 0.15 M NH4Cl

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Textbook Question

Determine whether or not the mixing of each pair of solutions results in a buffer. e. 125.0 mL of 0.15 M NH3; 150.0 mL of 0.20 M NaOH

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Textbook Question

Determine whether or not the mixing of each pair of solutions results in a buffer. c. 50.0 mL of 0.15 M HF; 20.0 mL of 0.15 M NaOH

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