Multiple Choice
Using the Henderson-Hasselbalch equation, what can you conclude about the relative concentrations of CH3NH2 and CH3NH3Cl in a buffer solution where the pH is 10.64 and the pKb of CH3NH2 is 3.36?
18. Aqueous Equilibrium
Henderson-Hasselbalch Equation
Multiple Choice
What mass of CH3COONa must be added to create 300.0 mL of a buffer solution with a pH of 5.00 and a CH3COOH concentration of 0.100 M? The pKa of acetic acid is 4.75.
A
1.23 g
B
0.82 g
C
3.69 g
D
2.46 g
Verified step by step guidance1
Identify the components of the buffer solution: acetic acid (CH3COOH) and sodium acetate (CH3COONa). The buffer solution is created by mixing these two components.
Use the Henderson-Hasselbalch equation to relate the pH of the buffer solution to the concentrations of the acid and its conjugate base: \( \text{pH} = \text{pKa} + \log \left( \frac{[\text{A}^-]}{[\text{HA}]} \right) \), where [A^-] is the concentration of the conjugate base (CH3COONa) and [HA] is the concentration of the acid (CH3COOH).
Substitute the given values into the Henderson-Hasselbalch equation: \( 5.00 = 4.75 + \log \left( \frac{[\text{CH}_3\text{COONa}]}{0.100} \right) \). Solve for the ratio \( \frac{[\text{CH}_3\text{COONa}]}{0.100} \).
Calculate the concentration of CH3COONa required using the ratio obtained from the previous step. Multiply the concentration by the volume of the solution (in liters) to find the number of moles of CH3COONa needed.
Convert the moles of CH3COONa to mass using its molar mass (82.03 g/mol). This will give you the mass of CH3COONa that must be added to the solution.
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