Question

A chemistry graduate student is given 450.mL of a 0.90M acetic acid HCH3CO2 solution. Acetic acid is a weak acid with Ka=1.8x10^−5.

What mass of KCH3CO2 should the student dissolve in the HCH3CO2 solution to turn it into a buffer with pH =4.35?

You may assume that the volume of the solution doesn't change when the KCH3CO2 is dissolved in it.

Be sure your answer has a unit symbol, and round it to 2 significant digits.

Answer 1

Henderson-Hasselbalch Equation is

pH = pKa + log ([KCH₃CO₂] / [HCH₃CO₂])

pKa = - log K_a

K_a = 1.8 X 10^-5

pKa = - log(1.8 X 10^-5) = 4.745

pH = 4.35

4.35 = 4.745 + log ([KCH₃CO₂] / [HCH₃CO₂])

log ([KCH₃CO₂] / [HCH₃CO₂]) = 4.35 - 4.745 = - 0.395

([KCH₃CO₂] / [HCH₃CO₂]) = 10^-0.395 = 0.403

Given [HCH₃CO₂] = 0.90 M

So [KCH₃CO₂] = 0.403 * [HCH₃CO₂] = 0.403 * 0.90 M = 0.36 M

Molarity = (mass of solute(g) * 1000) / (molar mass of solute(g/mol) * volume of solution(mL))

Molar mass of KCH₃CO₂ = 98.15 g/mol

Volume of solution = 450.0 mL

Mass of KCH₃CO₂ used = (0.36 M * 98.15 g/mol * 450.0 mL) / 1000 = 16 g