Multiple Choice
What is the pH of a solution when 40.00 mL of 0.1000 M HCl is added to a 50.00 mL aliquot of 0.1200 M ethylamine (CH3CH2NH2) with a Kb of 5.6 x 10^-4?
18. Aqueous Equilibrium
Titrations: Weak Base-Strong Acid
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Consider the titration of a 26.0 mL sample of 0.175 M CH3NH2 (Kb = 4.4 × 10⁻⁴) with 0.150 M HBr. Determine the pH at 5.0 mL of added acid.
A
9.24
B
10.18
C
10.52
D
11.02
Verified step by step guidance1
Start by identifying the type of titration: This is a weak base (CH3NH2) being titrated with a strong acid (HBr). The reaction will produce the conjugate acid (CH3NH3+) and water.
Calculate the initial moles of CH3NH2 using the formula: \( \text{moles} = \text{volume} \times \text{molarity} \). For CH3NH2, \( \text{moles} = 0.0260 \text{ L} \times 0.175 \text{ M} \).
Determine the moles of HBr added: \( \text{moles} = \text{volume} \times \text{molarity} \). For HBr, \( \text{moles} = 0.0050 \text{ L} \times 0.150 \text{ M} \).
Calculate the moles of CH3NH2 remaining after the reaction with HBr: Subtract the moles of HBr from the initial moles of CH3NH2. This will give you the moles of CH3NH2 left and the moles of CH3NH3+ formed.
Use the Henderson-Hasselbalch equation to find the pH: \( \text{pH} = \text{pKa} + \log \left( \frac{[\text{base}]}{[\text{acid}]} \right) \). First, find \( \text{pKa} \) from \( \text{pKb} \) using \( \text{pKa} + \text{pKb} = 14 \). Then, plug in the concentrations of CH3NH2 and CH3NH3+ to calculate the pH.
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