Methylamine (CH3NH2) is a weak base with a Kb = 4.38 x 10-4. What is the pH of a 0.295 M solution of methylamine?

Question

Methylamine (CH3NH2) is a weak base with a Kb = 4.38 x 10-4. What is the pH of a 0.295 M solution of methylamine?

Answer 1

Write the reaction of methylamine with water;

CH₃NH₂ + H₂O ==> CH₃NH₃⁺ + OH^- (the CH₃NH₂ is a base so receives H⁺ from H₂O, the acid)

Now, write the expression for the Kb:

Kb = [CH₃NH₃⁺] [OH^-] / [CH₃NH₂ ] and now fill in the known values and the unknowns:

4.38x10^-4 = (x)(x) / 0.295 - x (we will ignore the x in the denominator in order to avoid using the quadratic)

x2 = 1.29x10^-4

x = 1.14x10^-2 M = [OH^-] (NOTE: this is only ~3.9% of the 0.295 value above and so our assumption to omit it was valid according the "5% rule")

From the OH-, we can find pOH and the pH:

pOH = -log 1.14x10^-2 = 1.94

pH = 14 - 1.94

pH = 12.1