Multiple Choice
What is the concentration of H⁺ ions in a 2.0 M solution of acetic acid (HC₂H₃O₂) if the acid dissociation constant (Ka) is 1.8 x 10⁻⁵?
17. Acid and Base Equilibrium
pH of Weak Acids
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What is the pH of a 0.200 M CH3NH3Br solution, given that the Kb of CH3NH2 is 4.4 × 10⁻⁴?
A
3.72
B
7.00
C
5.28
D
8.72
Verified step by step guidance1
Identify that CH3NH3Br is a salt formed from CH3NH2 (a weak base) and HBr (a strong acid). In solution, CH3NH3+ acts as a weak acid.
Write the hydrolysis reaction for CH3NH3+: CH3NH3+ + H2O ⇌ CH3NH2 + H3O+. This reaction shows how CH3NH3+ can donate a proton to water, forming CH3NH2 and H3O+.
Use the relationship between Ka and Kb for conjugate acid-base pairs: Ka × Kb = Kw, where Kw is the ion-product constant of water (1.0 × 10⁻¹⁴ at 25°C). Calculate Ka for CH3NH3+ using the given Kb for CH3NH2.
Set up an ICE (Initial, Change, Equilibrium) table for the hydrolysis reaction to find the concentration of H3O+ at equilibrium. Assume the initial concentration of CH3NH3+ is 0.200 M and the initial concentration of H3O+ is approximately 0.
Use the expression for Ka: Ka = [CH3NH2][H3O+]/[CH3NH3+]. Substitute the equilibrium concentrations from the ICE table into this expression to solve for [H3O+]. Finally, calculate the pH using the formula pH = -log[H3O+].
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