chemistry graduate student is given 450. mL of a 1.10M nitrous acid (HNO,) solution. Nitrous acid is a weak acid with = 4.5 × 10 ". What mass of NaNO, should the student dissolve in the HNO, solution to turn it into a buffer with = 3.18? u may assume that the volume of the solution doesn't change when the NaNO, is dissolved in it. Be sure your answer has unit symbol, and round it to 2 significant digits. Ox10 O
chemistry graduate student is given 450. mL of a 1.10M nitrous acid (HNO,) solution. Nitrous acid is a weak acid with = 4.5 × 10 ". What mass of NaNO, should the student dissolve in the HNO, solution to turn it into a buffer with = 3.18? u may assume that the volume of the solution doesn't change when the NaNO, is dissolved in it. Be sure your answer has unit symbol, and round it to 2 significant digits. Ox10 O
Introductory Chemistry: A Foundation
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Chapter16: Acids And Bases
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Problem 65AP: . The concepts of acid-base equilibria were developed in this chapter for aqueous solutions (in...
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![### Problem Statement: Buffer Solution Preparation
**Objective:**
Calculate the mass of sodium nitrite (NaNO₂) that should be dissolved in a nitrous acid (HNO₂) solution to achieve a desired pH.
**Given Data:**
- Volume of HNO₂ solution: 450 mL
- Initial concentration of HNO₂ solution: 1.10 M
- Acid dissociation constant (Kₐ) for HNO₂: 4.5 × 10⁻⁴
- Desired pH of the buffer solution: 3.18
**Assumptions:**
- The volume of the solution does not change when NaNO₂ is dissolved.
**Steps for Calculation:**
1. **Henderson-Hasselbalch Equation:**
To find the required concentration of the conjugate base (NO₂⁻), we use the Henderson-Hasselbalch equation:
\[
\text{pH} = \text{p}K_a + \log \left( \frac{[\text{A}^-]}{[\text{HA}]} \right)
\]
where:
- \(\text{pH} = 3.18\)
- \(\text{p}K_a = - \log (4.5 \times 10^{-4}) = 3.35\)
2. **Rearrange the Equation:**
Solving for the ratio of conjugate base (NO₂⁻) to acid (HNO₂):
\[
3.18 = 3.35 + \log \left( \frac{[NO_2^-]}{[HNO_2]} \right)
\]
\[
\log \left( \frac{[NO_2^-]}{[HNO_2]} \right) = 3.18 - 3.35 = -0.17
\]
\[
\frac{[NO_2^-]}{[HNO_2]} = 10^{-0.17} \approx 0.68
\]
3. **Find the Concentration of NO₂⁻:**
\[
[NO_2^-] = 0.68 \times [HNO_2]
\]
Since the initial concentration of HNO₂ is 1.](/v2/_next/image?url=https%3A%2F%2Fcontent.bartleby.com%2Fqna-images%2Fquestion%2F38d49af3-6887-43af-8a6c-6e39b942bd2d%2F03c57f78-db24-4b8f-975a-fd2cb8e04877%2Fpvb8u5d_processed.png&w=3840&q=75)
Transcribed Image Text:### Problem Statement: Buffer Solution Preparation
**Objective:**
Calculate the mass of sodium nitrite (NaNO₂) that should be dissolved in a nitrous acid (HNO₂) solution to achieve a desired pH.
**Given Data:**
- Volume of HNO₂ solution: 450 mL
- Initial concentration of HNO₂ solution: 1.10 M
- Acid dissociation constant (Kₐ) for HNO₂: 4.5 × 10⁻⁴
- Desired pH of the buffer solution: 3.18
**Assumptions:**
- The volume of the solution does not change when NaNO₂ is dissolved.
**Steps for Calculation:**
1. **Henderson-Hasselbalch Equation:**
To find the required concentration of the conjugate base (NO₂⁻), we use the Henderson-Hasselbalch equation:
\[
\text{pH} = \text{p}K_a + \log \left( \frac{[\text{A}^-]}{[\text{HA}]} \right)
\]
where:
- \(\text{pH} = 3.18\)
- \(\text{p}K_a = - \log (4.5 \times 10^{-4}) = 3.35\)
2. **Rearrange the Equation:**
Solving for the ratio of conjugate base (NO₂⁻) to acid (HNO₂):
\[
3.18 = 3.35 + \log \left( \frac{[NO_2^-]}{[HNO_2]} \right)
\]
\[
\log \left( \frac{[NO_2^-]}{[HNO_2]} \right) = 3.18 - 3.35 = -0.17
\]
\[
\frac{[NO_2^-]}{[HNO_2]} = 10^{-0.17} \approx 0.68
\]
3. **Find the Concentration of NO₂⁻:**
\[
[NO_2^-] = 0.68 \times [HNO_2]
\]
Since the initial concentration of HNO₂ is 1.
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