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Ch.17 - Aqueous Ionic Equilibrium
All textbooksTro 4th EditionCh.17 - Aqueous Ionic EquilibriumProblem 49
Chapter 17, Problem 49

For each solution, calculate the initial and final pH after adding 0.010 mol of HCl: a. 500.0 mL of pure water b. 500.0 mL of a buffer solution that is 0.125 M in HC2H3O2 and 0.115 M in NaC2H3O2 c. 500.0 mL of a buffer solution that is 0.155 M in C2H5NH2 and 0.145 M in C2H5NH3Cl.

Verified step by step guidance
1
Identify the type of solution for each part: pure water, acetic acid/acetate buffer, and ethylamine/ethylammonium chloride buffer.
For part (a), calculate the initial pH of pure water, which is neutral at pH 7. Then, determine the change in pH after adding 0.010 mol of HCl by calculating the concentration of HCl in the solution and using the formula for pH: \( \text{pH} = -\log[H^+] \).
For part (b), use the Henderson-Hasselbalch equation to calculate the initial pH of the acetic acid/acetate buffer: \( \text{pH} = \text{pK}_a + \log \left( \frac{[\text{A}^-]}{[\text{HA}]} \right) \), where \( \text{pK}_a \) is the acid dissociation constant for acetic acid. Then, calculate the final pH after adding HCl by adjusting the concentrations of acetic acid and acetate ion and applying the Henderson-Hasselbalch equation again.
For part (c), calculate the initial pH of the ethylamine/ethylammonium chloride buffer using the Henderson-Hasselbalch equation: \( \text{pH} = \text{pK}_b + \log \left( \frac{[\text{B}]}{[\text{BH}^+]} \right) \), where \( \text{pK}_b \) is the base dissociation constant for ethylamine. After adding HCl, adjust the concentrations of ethylamine and ethylammonium ion and use the Henderson-Hasselbalch equation to find the final pH.
Summarize the effect of adding a strong acid (HCl) to each solution, highlighting the buffering capacity of the buffer solutions compared to pure water.>

Key Concepts

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

pH Scale

The pH scale measures the acidity or basicity of a solution, ranging from 0 to 14. A pH of 7 is neutral, below 7 indicates acidity, and above 7 indicates basicity. The pH is calculated using the formula pH = -log[H+], where [H+] is the concentration of hydrogen ions in the solution. Understanding pH is crucial for predicting how the addition of acids or bases will affect a solution.
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Buffer Solutions

Buffer solutions are mixtures that resist changes in pH when small amounts of acid or base are added. They typically consist of a weak acid and its conjugate base or a weak base and its conjugate acid. The Henderson-Hasselbalch equation can be used to calculate the pH of buffer solutions, which is essential for understanding how they maintain stability in pH during the addition of HCl in the given scenarios.
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Dilution and Concentration Changes

When a solute is added to a solution, the concentration of the solute changes, which can affect the pH. In the case of adding HCl to water or buffer solutions, the initial concentration of H+ ions must be considered, along with the volume of the solution. The final pH can be calculated by determining the new concentration of H+ ions after the addition of HCl, which is critical for accurately predicting the pH changes in the solutions.
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Related Practice
Textbook Question

A 100.0-mL buffer solution is 0.175 M in HClO and 0.150 M in NaClO. a. What is the initial pH of this solution?

Textbook Question

A 100.0-mL buffer solution is 0.175 M in HClO and 0.150 M in NaClO. b. What is the pH after addition of 150.0 mg of HBr?

Textbook Question

A 100.0-mL buffer solution is 0.175 M in HClO and 0.150 M in NaClO. c. What is the pH after addition of 85.0 mg of NaOH?

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

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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