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I did a lab where I tested the solubility of carbon dioxide in water at different temperatures. I took 50 mL of carbonated water at different temperatures and then I titrated that with sodium hydroxide solution. The indicator I used was phenolphthalein with an endpoint of 8.3 pH. I'm unsure how to calculate the amount of carbon dioxide that was dissolved. I'm stuck about deciding which equation to use.

This equation would mean that there is a 1:1 ratio. $$\ce{CO2(aq) + NaOH(aq) → NaHCO3(aq)}$$ And this equation would mean that there is a 1:2 ratio.

$$\ce{CO2(aq) + H2O(l)->H2CO3(aq)}$$ $$\ce{H2CO3(aq) + 2NaOH(aq)->2H2O(l) + Na2CO3(aq)}$$

Which one would be the correct equation to use?

Also, there are literature values for the solubility of carbon dioxide in water at different temperatures. Do those solubility values apply to my lab? I ask this because I measured the amount of carbon dioxide present in a given amount of carbonated water, I didn't add carbon dioxide to pure water to measure how much can actually dissolve.

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  • $\begingroup$ What was the endpoint pH of your titration? How were you monitoring it? Colorimetrically? If so, with what indicator? Or did you monitor with a pH probe? $\endgroup$ – Curt F. Mar 6 '16 at 17:40
  • $\begingroup$ I used phenolphthalein, the end point was around 8.3. $\endgroup$ – Yulmart Mar 6 '16 at 17:41
  • $\begingroup$ Great -- could you edit the question to include that detail? It's very important in determining the right answer, as you might infer from A.K.'s answer. $\endgroup$ – Curt F. Mar 6 '16 at 17:42
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Sodium carbonate ($\ce{NaCO3}$) is basic and gives solutions with a pH of around 10. Sodium bicarbonate ($\ce{NaHCO3}$) creates solutions with a pH of approximately 7.3. This is in fact what the human body uses to regulate pH. You want to used the sodium bicarbonate since it will yield a nearly neutral solution.

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Compare your endpoint $\ce{OH}$ with the relevant $pK_a$ values. From Wikpedia:

$$\begin{align}\\ \ce{CO2 + H2O = HCO3- + H+} & & {pK_\mathrm{a}=6.3}\ \ \\ \ce{HCO3- = CO3^2- + H+} & & {pK_\mathrm{a}=10.3}\\ \end{align}$$

The first of these equilibria is for $\ce{CO2}$ as the solute. Remember that a base will deprotonate a weak acid only as the pH surpasses the $pK_\mathrm{a}$ for that acid. So an endpoint pH of $8.3$ allows only the bicarbonate to form.

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