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Ch.18 - Aqueous Ionic Equilibrium
Tro - Chemistry: A Molecular Approach 6th Edition
Tro6th EditionChemistry: A Molecular ApproachISBN: 9780137832217Not the one you use?Change textbook
All textbooksTro 6th EditionCh.18 - Aqueous Ionic EquilibriumProblem 55c
Chapter 18, Problem 55c

Determine whether or not the mixing of each pair of solutions results in a buffer. c. 55.0 mL of 0.15 M HF; 85.0 mL of 0.10 M NaF

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1
Identify the components of the solutions: HF is a weak acid and NaF is its conjugate base.
Calculate the moles of HF using the formula: moles = concentration \( \times \) volume. Convert the volume from mL to L before calculation.
Calculate the moles of NaF using the same formula: moles = concentration \( \times \) volume. Again, convert the volume from mL to L before calculation.
Determine if the solution can act as a buffer: A buffer solution contains a weak acid and its conjugate base in significant amounts.
Compare the moles of HF and NaF to ensure both are present in significant amounts to resist changes in pH, confirming the presence of a buffer.

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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. In this case, HF (hydrofluoric acid) is a weak acid, and NaF (sodium fluoride) provides the conjugate base, F-. Together, they can form a buffer solution.
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Buffer Solutions

Weak Acids and Conjugate Bases

Weak acids are substances that partially dissociate in solution, establishing an equilibrium between the undissociated acid and its ions. The conjugate base is formed when the weak acid donates a proton (H+). In the given example, HF partially dissociates to produce H+ and F-, making NaF a source of F-, which can react with added H+ to maintain pH stability.

Henderson-Hasselbalch Equation

The Henderson-Hasselbalch equation is a mathematical formula used to calculate the pH of a buffer solution. It is expressed as pH = pKa + log([A-]/[HA]), where [A-] is the concentration of the conjugate base and [HA] is the concentration of the weak acid. This equation helps determine if the given concentrations of HF and NaF will effectively create a buffer by assessing their ratio.
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Related Practice
Textbook Question

Determine whether or not the mixing of each pair of solutions results in a buffer. b.155.0 mL of 0.15 M NH3 ; 155.0 mL of 0.10 M HCl

Textbook Question

Determine whether or not the mixing of each pair of solutions results in a buffer. d. 105.0 mL of 0.12 M CH3NH2 ; 110.0 mL of 0.15 M CH3NH3Cl

Textbook Question

Determine whether or not the mixing of each pair of solutions results in a buffer. c. 225.0 mL of 0.10 M NH3 ; 250.0 mL of 0.15 M NH4Cl

Textbook Question

A 100.0-mL buffer solution is 0.100 M in NH3 and 0.125 M in NH4Br. What mass of HCl can this buffer neutralize before the pH falls below 9.00?

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Textbook Question

For each solution, calculate the initial and final pH after adding 0.010 mol of NaOH. a. 250.0 mL of pure water b. 250.0 mL of a buffer solution that is 0.195 M in HCHO2 and 0.275 M in KCHO2 c. 250.0 mL of a buffer solution that is 0.255 M in CH3CH2NH2 and 0.235 M in CH3CH2NH3Cl

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Textbook Question

Determine whether or not the mixing of each pair of solutions results in a buffer. a. 155.0 mL of 0.15 M NH3 ; 175.0 mL of 0.17 M HCl