Multiple Choice
Calculate the pH of a solution that is 0.540 M in pyridine (C5H5N) and 0.470 M in pyridinium chloride (C5H5NHCl). The Kb for pyridine is 1.7 x 10^-9.
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17. Acid and Base Equilibrium
pH of Weak Bases
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?
A
9.85
B
7.00
C
11.22
D
2.78
Verified step by step guidance1
First, identify that CH3NH2 is a weak base. To find the initial pH, we need to calculate the concentration of OH⁻ ions in the solution.
Use the base dissociation constant (Kb) for CH3NH2, which is given as 4.4 × 10⁻⁴. The equilibrium expression for the dissociation of CH3NH2 in water is:
Set up an ICE table (Initial, Change, Equilibrium) for the dissociation of CH3NH2. Initially, the concentration of CH3NH2 is 0.185 M, and the concentrations of CH3NH3⁺ and OH⁻ are 0 M.
Assume that x is the change in concentration for CH3NH3⁺ and OH⁻ at equilibrium. Therefore, the equilibrium concentrations are:
,
,
Substitute these equilibrium concentrations into the Kb expression and solve for x, which represents [OH⁻]. Then, calculate the pOH using the formula:
. Finally, convert pOH to pH using:
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