Multiple Choice
Consider the titration of a 26.0-mL sample of 0.185 M CH3NH2 (Kb = 4.4 × 10⁻⁴) with 0.155 M HBr. What is the initial pH of the CH3NH2 solution before any HBr is added?
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17. Acid and Base Equilibrium
pH of Weak Bases
Multiple Choice
For a solution that is 0.164 M NH3 and 0.110 M NH4Cl, calculate the concentration of [H3O+].
A
5.6 x 10^-9 M
B
7.4 x 10^-8 M
C
1.0 x 10^-5 M
D
3.2 x 10^-7 M
Verified step by step guidance1
Identify that the solution is a buffer solution composed of a weak base (NH3) and its conjugate acid (NH4Cl). This means you can use the Henderson-Hasselbalch equation to find the pH.
Write the Henderson-Hasselbalch equation for a basic buffer: \( \text{pH} = \text{pK}_a + \log \left( \frac{[\text{base}]}{[\text{acid}]} \right) \). Here, the base is NH3 and the acid is NH4Cl.
Find the \( \text{pK}_a \) of NH4+, which is the conjugate acid of NH3. You can use the relation \( \text{pK}_a = 14 - \text{pK}_b \), where \( \text{pK}_b \) is the base dissociation constant of NH3.
Substitute the concentrations of NH3 (0.164 M) and NH4Cl (0.110 M) into the Henderson-Hasselbalch equation along with the calculated \( \text{pK}_a \) to find the pH of the solution.
Convert the pH to \([\text{H}_3\text{O}^+]\) concentration using the formula \([\text{H}_3\text{O}^+] = 10^{-\text{pH}}\). This will give you the concentration of hydronium ions in the solution.
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