How many grams of solid sodium nitrite should be added to 2.00 L of 0.117 M nitrous acid solution to prepare a buffer with a pH of 4.250? (K₂ for nitrous acid = 4.50E-4) g sodium nitrite
How many grams of solid sodium nitrite should be added to 2.00 L of 0.117 M nitrous acid solution to prepare a buffer with a pH of 4.250? (K₂ for nitrous acid = 4.50E-4) g sodium nitrite
Chemistry: Principles and Practice
3rd Edition
ISBN:9780534420123
Author:Daniel L. Reger, Scott R. Goode, David W. Ball, Edward Mercer
Publisher:Daniel L. Reger, Scott R. Goode, David W. Ball, Edward Mercer
Chapter16: Reactions Between Acids And Bases
Section: Chapter Questions
Problem 16.7QE
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![**Buffer Solution Preparation Calculation**
**Problem Statement:**
How many grams of solid sodium nitrite should be added to 2.00 L of 0.117 M nitrous acid solution to prepare a buffer with a pH of 4.250? ( \( K_a \) for nitrous acid = \( 4.50 \times 10^{-4} \) )
**Solution:**
To solve this problem, we will use the Henderson-Hasselbalch equation for buffer solutions:
\[ \text{pH} = \text{p}K_a + \log \left( \frac{[\text{A}^-]}{[\text{HA}]} \right) \]
Where:
- \( \text{pH} \) is the desired pH of the buffer solution.
- \( \text{p}K_a \) is the negative logarithm of the acid dissociation constant (\( K_a \)).
- \([\text{A}^-]\) is the concentration of the conjugate base.
- \([\text{HA}]\) is the concentration of the weak acid.
1. **Calculate \( \text{p}K_a \):**
\[ \text{p}K_a = -\log ( 4.50 \times 10^{-4} ) \]
2. **Set up the Henderson-Hasselbalch equation and solve for \([\text{A}^-]\):**
\[ 4.250 = \text{p}K_a + \log \left( \frac{[\text{A}^-]}{0.117} \right) \]
3. **Determine the concentration of \([\text{A}^-]\) which is the concentration of sodium nitrite.**
4. **Calculate the moles of sodium nitrite required:**
\[ \text{Moles of sodium nitrite} = \text{Concentration} \times \text{Volume} \]
5. **Convert moles to grams:**
\[ \text{Grams of sodium nitrite} = \text{Moles} \times \text{Molar Mass} \]
Insert the numbers accordingly and solve for the mass of sodium nitrite needed.
**Input Field:**
\[ \boxed{ \phantom{000} } \, \text{g sodium nitrite} \]
**Note:**
This explanation does not include a graph or diagram](/v2/_next/image?url=https%3A%2F%2Fcontent.bartleby.com%2Fqna-images%2Fquestion%2F6ab0bbaf-7c6a-41d9-a28d-d2a0ce6f8c93%2F3b8a2556-ec76-4223-adab-f1125b88b048%2Fomh2ba9_processed.png&w=3840&q=75)
Transcribed Image Text:**Buffer Solution Preparation Calculation**
**Problem Statement:**
How many grams of solid sodium nitrite should be added to 2.00 L of 0.117 M nitrous acid solution to prepare a buffer with a pH of 4.250? ( \( K_a \) for nitrous acid = \( 4.50 \times 10^{-4} \) )
**Solution:**
To solve this problem, we will use the Henderson-Hasselbalch equation for buffer solutions:
\[ \text{pH} = \text{p}K_a + \log \left( \frac{[\text{A}^-]}{[\text{HA}]} \right) \]
Where:
- \( \text{pH} \) is the desired pH of the buffer solution.
- \( \text{p}K_a \) is the negative logarithm of the acid dissociation constant (\( K_a \)).
- \([\text{A}^-]\) is the concentration of the conjugate base.
- \([\text{HA}]\) is the concentration of the weak acid.
1. **Calculate \( \text{p}K_a \):**
\[ \text{p}K_a = -\log ( 4.50 \times 10^{-4} ) \]
2. **Set up the Henderson-Hasselbalch equation and solve for \([\text{A}^-]\):**
\[ 4.250 = \text{p}K_a + \log \left( \frac{[\text{A}^-]}{0.117} \right) \]
3. **Determine the concentration of \([\text{A}^-]\) which is the concentration of sodium nitrite.**
4. **Calculate the moles of sodium nitrite required:**
\[ \text{Moles of sodium nitrite} = \text{Concentration} \times \text{Volume} \]
5. **Convert moles to grams:**
\[ \text{Grams of sodium nitrite} = \text{Moles} \times \text{Molar Mass} \]
Insert the numbers accordingly and solve for the mass of sodium nitrite needed.
**Input Field:**
\[ \boxed{ \phantom{000} } \, \text{g sodium nitrite} \]
**Note:**
This explanation does not include a graph or diagram
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