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 pH after adding 6.0 mL of HBr?
18. Aqueous Equilibrium
Titrations: Weak Base-Strong Acid
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Pyridine, C5H5N, is a weak base with Kb = 1.5 x 10^-9. What is the pH at the equivalence point when 238 mL of 0.78 M pyridine is titrated with 0.78 M HCl?
A
11.00
B
7.00
C
4.98
D
9.02
Verified step by step guidance1
Identify that pyridine (C5H5N) is a weak base and is being titrated with a strong acid (HCl). At the equivalence point, all the pyridine will be converted to its conjugate acid, pyridinium ion (C5H5NH+).
Calculate the moles of pyridine initially present using the formula: \( \text{moles} = \text{concentration} \times \text{volume} \). Use the given concentration (0.78 M) and volume (238 mL, converted to liters).
At the equivalence point, the moles of pyridine will equal the moles of HCl added. Therefore, the concentration of pyridinium ion (C5H5NH+) in the solution can be calculated using the total volume of the solution after titration.
Recognize that the pyridinium ion (C5H5NH+) is a weak acid. Use the relationship between the base dissociation constant (Kb) and the acid dissociation constant (Ka) for the conjugate acid: \( K_a = \frac{K_w}{K_b} \), where \( K_w = 1.0 \times 10^{-14} \).
Use the expression for the ionization of the weak acid (C5H5NH+) to find the concentration of hydrogen ions \( [H^+] \) and then calculate the pH using the formula: \( \text{pH} = -\log[H^+] \).
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