Multiple Choice
What is the pH of a 0.16 M NH4Cl solution, given that the Kb for NH3 is 1.76 × 10⁻⁵?
17. Acid and Base Equilibrium
pH of Weak Acids
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Determine the pH of a buffer solution prepared by dissolving 0.20 mol of formic acid (HCHO) and 0.80 mol of sodium formate (NaCHO) in enough water to make 1.0 L of solution. The pKa of formic acid is 3.75.
A
3.50
B
3.75
C
4.00
D
4.25
Verified step by step guidance1
Identify the components of the buffer solution: formic acid (HCHO) and sodium formate (NaCHO). Formic acid is the weak acid, and sodium formate is its conjugate base.
Use the Henderson-Hasselbalch equation to calculate the pH of the buffer solution: \( \text{pH} = \text{pKa} + \log \left( \frac{[\text{A}^-]}{[\text{HA}]} \right) \). Here, \([\text{A}^-]\) is the concentration of the conjugate base (sodium formate), and \([\text{HA}]\) is the concentration of the weak acid (formic acid).
Calculate the concentrations of the acid and base in the solution. Since the solution volume is 1.0 L, the concentration of formic acid \([\text{HA}]\) is 0.20 M, and the concentration of sodium formate \([\text{A}^-]\) is 0.80 M.
Substitute the given values into the Henderson-Hasselbalch equation: \( \text{pH} = 3.75 + \log \left( \frac{0.80}{0.20} \right) \).
Evaluate the logarithmic expression \( \log \left( \frac{0.80}{0.20} \right) \) and add it to the pKa value to find the pH of the buffer solution.
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