**Buffer Solution Calculation Example**A buffer solution is prepared with the following concentrations:- 0.372 M (Molar) of NaHC₂O₄- 0.323 M (Molar) of Na₂C₂O₄Given the acid dissociation constant \( K_a \) for HC₂O₄⁻ is \( 6.4 \times 10^{-5} \).**Question:** What is the pH of this buffer solution?This exercise involves the use of the Henderson-Hasselbalch equation for calculating the pH of a buffer solution:\[ \text{pH} = pK_a + \log \left( \frac{[\text{Conjugate Base}]}{[\text{Weak Acid}]} \right) \]**Explanation:**1. \( K_a \) is provided, and from \( K_a \), \( pK_a \) can be found using the equation: \[ pK_a = -\log(K_a) \]2. Substitute the concentrations of the conjugate base \( [\text{Na}_{2}\text{C}_2\text{O}_{4}] \) and the weak acid \( [\text{NaHC}_2\text{O}_{4}] \) into the Henderson-Hasselbalch equation.The given information is used to find the pH of the buffer through proper calculation steps involving logarithms and the provided constants. Use this example to understand the practical application of acid-base chemistry in buffer systems and the importance of accurate calculations in determining pH levels.

Given: Concentration of NaHC2O4 is 0.372 M Concentration of Na2HC2O4 is 0.372 M Ka is 6.4×10-5 To…

Answered: A buffer solution is 0.372 M in NaHC204…

A buffer solution is 0.372 M in NaHC204 and 0.323 M in Na2C₂O4. If K₂ for HC204 is 6.4E-5, what is the pH of this buffer solution?

Chemistry

10th Edition

ISBN:9781305957404

Author:Steven S. Zumdahl, Susan A. Zumdahl, Donald J. DeCoste

Publisher:Steven S. Zumdahl, Susan A. Zumdahl, Donald J. DeCoste

Chapter1: Chemical Foundations

Section: Chapter Questions

Problem 1RQ: Define and explain the differences between the following terms. a. law and theory b. theory and...

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Concept explainers

Acid Base Titration

An acid-base titration is the method of quantitative analysis for determining the concentration of an acid and base by exactly neutralizing it with a standard solution of base or acid having known concentration.

Question

**Buffer Solution Calculation Example**

A buffer solution is prepared with the following concentrations:
- 0.372 M (Molar) of NaHC₂O₄
- 0.323 M (Molar) of Na₂C₂O₄

Given the acid dissociation constant \( K_a \) for HC₂O₄⁻ is \( 6.4 \times 10^{-5} \).

**Question:** What is the pH of this buffer solution?

This exercise involves the use of the Henderson-Hasselbalch equation for calculating the pH of a buffer solution:
\[ \text{pH} = pK_a + \log \left( \frac{[\text{Conjugate Base}]}{[\text{Weak Acid}]} \right) \]

**Explanation:**

1. \( K_a \) is provided, and from \( K_a \), \( pK_a \) can be found using the equation:
\[ pK_a = -\log(K_a) \]

2. Substitute the concentrations of the conjugate base \( [\text{Na}_{2}\text{C}_2\text{O}_{4}] \) and the weak acid \( [\text{NaHC}_2\text{O}_{4}] \) into the Henderson-Hasselbalch equation.

The given information is used to find the pH of the buffer through proper calculation steps involving logarithms and the provided constants.

Use this example to understand the practical application of acid-base chemistry in buffer systems and the importance of accurate calculations in determining pH levels.

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