[jazz519]

Hi, for the question below I'm having trouble. I thought of two approaches but both give me different answers

Calculate the pH of a solution that is made up to contain an analytical concentration of 0.03 M in HCO3^- and 0.06 M in CO3^2-.

The base dissociation constants (Kb) for CO3^2- and HCO3^- are 2.13 x 10^-4 and 2.40 x 10^-8   respectively.

For the approach to this problem I'm unsure if the way I should do it is through two separate ICE table calculations for OH^- using the reactions of carbonate to hydrogen carbonate and then hydrogen carbonate to carbonic acid, and then add those two OH^- values together to give the total OH^-, which I could then use to find pH.

Or if can just use the Henderson-Hasselbalch equation and do pOH = pKb(co3^2-) + log10([hco3^-]/[co3^2-)] and then calculate pH from that, but my problem with this second approach doesn't it ignore that the hco3^- can also act as a base in another equilibrium or can we just use this approach because the pKb values are further than 3 apart

Any guidance would be helpful as to which method to use, thank you

[AWK]

Simply use the Henderson-Hasselbach equation. The errors caused by the omission of other equilibria in the solution will be much smaller than the error caused by the omission of the influence of ionic strength (I guess - over 0.1 pH unit), which you do not have to account for.

[jazz519]

Oh okay makes sense thank you

[Borek]

This is a bit ambiguous but the way I read the question there are no "other equilibria" that can change anything in the solution. You are told that the solution contains CO₃^2- and HCO₃^- at given concentrations - so these are equilibrium concentrations that won't change in time. That in turn means that HH equation (which is nothing else but just a rearranged dissociation constant) must hold.

If you were told "solution was prepared by putting ions at concentrations" that would mean they were allowed to react further.

[jazz519]

That's a good point and yeah I was confused by whether that was the case but thanks for clearing it up