Multiple Choice
Consider the titration of a 25.0-mL sample of 0.175 M CH3NH2 with 0.150 M HBr. What is the initial pH of the CH3NH2 solution before any HBr is added?
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 the equivalence point.
A
pH = 5.28
B
pH = 9.25
C
pH = 7.00
D
pH = 3.75
Verified step by step guidance1
Identify the reaction: CH₃NH₂ (a weak base) reacts with HBr (a strong acid) to form CH₃NH₃⁺ and Br⁻. The reaction is: CH₃NH₂ + HBr → CH₃NH₃⁺ + Br⁻.
Determine the moles of CH₃NH₂ initially present using the formula: moles = concentration × volume. Convert the volume from mL to L before calculating.
At the equivalence point, the moles of HBr added will equal the moles of CH₃NH₂ initially present. Calculate the volume of HBr needed using its concentration.
At the equivalence point, the solution contains CH₃NH₃⁺, which is the conjugate acid of CH₃NH₂. Use the Kb of CH₃NH₂ to find the Ka of CH₃NH₃⁺ using the relation: K_w = K_a × K_b, where K_w is the ion-product constant of water (1.0 × 10⁻¹⁴ at 25°C).
Calculate the pH of the solution at the equivalence point by setting up an equilibrium expression for the dissociation of CH₃NH₃⁺ in water: CH₃NH₃⁺ ⇌ CH₃NH₂ + H⁺. Use the calculated Ka to find [H⁺], and then determine the pH using the formula: pH = -log[H⁺].
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