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?
17. Acid and Base Equilibrium
pH of Weak Bases
Struggling with General Chemistry?
Join thousands of students who trust us to help them ace their exams!Watch the first videoMultiple 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.
0:51mWatch next
Master ICE Charts of Weak Bases with a bite sized video explanation from Jules
Start learningRelated Videos
Related Practice
pH of Weak Bases practice set
Problem sets built by lead tutorsExpert video explanations