I am having trouble with this chemistry problem

Question

A buffer solution is made such that the initial concentrations of lactic acid (HC3H5O3) and the lactate ion (C3HO) are .6M and .620M, respectively. What is the resulting pH if 100.0 mL of .2M potassium hydroxide is added to .300L of the buffer solution? (The Ka of HC3H5O3 is 1.4e-4)

Answer 1

Henderson-Hasselbalch:

pH = pKa + log([A^-]/[HA] where pKa is -log(Ka)

Calculate moles of A^-: before adding OH^- (.62 mol/l)(.3l) = .186 moles

Calculate moles of HA before adding OH^- (.6 mol/l)(.3 l) = .18 moles

Calculate the )H^- introduced (.2 mol/l)(.1 l) = .02 moles

The lactate increases by .02 moles and the lactic acid decreases by .02 moles

[A^-] = .206 and [HA] = .18

Plug into equation to find pH. You can use moles instead of molarity as the Bronsted-Lowry pair are 1 to 1 and the total volume is the same for both the acid and conjugate base.

Strictly speaking, you should analyze this using ICE table, but for weak acid - strong base titration far from the equivalence point, you can assume that the strong bases dominates.

I hope that helped.