Hi, Id like to ask some opinions regarding a problem I have. Im normally pretty good with these problems, but this one bugs me.
Calculate the pH of a solution obtained by mixing 100mL of H2S 0.1M with a) 200mL KOH 0.05M; b) 300ml KOH 0.05M
The constants of dissociation are given: Ka1 = 1.1 x 10^-7 and Ka2 = 10^-14
For a) I first note that the two compounds are equimolar, so the reaction H2S + KOH --> HS- + H20 + K+ proceeds untill the total consumption of both compounds, after which I have a solution of the conjugate base, for which I can calculate the pH by means of [OH] = (Kb x Cb)^1/2, after finding Kb with Kb = Kw / Ka1 and adjusting the concentration to the new volume.
My real problem lies in part b).
I now have a excess of KOH in respect to H2S. Normally I would argue that the reaction HS- + KOH --> H20 + S2- + K+ would then proceed towarss the consumption of KOH. If this is the case then I wouldnt know how to get the pH because in addition to the equilibrium HS- / S2-, I also have the an equilibrium with H2S / HS-.