I've conducted an experiment where a solution of $\pu{4.2 M}$ $\ce{NaOH}$ was left exposed to atmospheric $\ce{CO2}$ for a period of time and titrated using potentiometric titration and potassium hydrogen phthalate ($\ce{KHP}$) to investigate the degradation of $\ce{NaOH}$ over time. I was under the impression that the reaction for such a high pH solution would look like this:
$$\ce{2NaOH(aq) + CO2 <=> Na2CO3(aq) + H2O} $$
But the results I have got don't fit this. For example the rate of change of $[\ce{NaOH}]$ and $[\ce{Na2CO3}]$ should be a 2:1 ratio but I got 0.849:1. This is a significant difference. The data I have seem accurate and the titration curves lined up well on each run, I don't think this is a significant error on my part so I'm now wondering if instead of the above reaction, $\ce{NaHCO3}$ was being formed:
$$\ce{NaOH(aq) + CO2(g) <=> NaHCO3(aq)}$$
This would fit the ratio I got much better but I thought this wasn't possible at a pH of 10 or above? Here are my titration and Concentration vs time graphs. I had previously reasoned that the fact I was getting only two equivalence points was due to the third step being a weak neutralization with a pH of about 7 that would barely show but now it seems like it could be because $\ce{NaHCO3}$ is the reacting not $\ce{Na2CO3}$.
Any help working out what is happening here would be great...