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The pH at the equivalence point of titration of $\ce{Na2CO3}$ solution with $\ce{HCl}$ is around 3.7, as shown in this titration curve:

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At this point, the following reaction is completed:
$$\ce{NaHCO3 + HCl -> NaCl + H2O + CO2}$$ But considering the fact that $\ce{NaCl}$ is a neutral salt, shouldn't the pH be closer to 7? The $\ce{H2CO3}$ formed also decomposes to form $\ce{H2O + CO2}$, and shouldn't affect the pH either. Why then is the equivalence point at such a low pH value?

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  • $\begingroup$ Contrary to what you say, CO2 in water does affect the pH. You need a little bit more HCl to obtain the equivalence point. $\endgroup$
    – Maurice
    Commented Nov 27, 2020 at 14:34
  • $\begingroup$ @Maurice, Wouldn't the equivalence point be the point where the solution has been completely neutralized, and thus would only contain $\ce{NaCl + H2CO3}$? Also doesn't the $\ce{CO2}$ formed escape, and not cause a significant decrease in pH? $\endgroup$
    – jerrythealpaca
    Commented Nov 27, 2020 at 15:40
  • $\begingroup$ The equivalence point is not the point where the solution contains $\ce{NaCl}$ and that's all. No. The equivalent point is when the solution contains $\ce{NaCl}$ plus the small amount of $\ce{CO2}$ which remains dissolved. And this amount makes the solution a bit acidic. So the pH at the equivalence point is not $7$, but a bit less. $\endgroup$
    – Maurice
    Commented Nov 27, 2020 at 17:25
  • $\begingroup$ @Maurice, You are mixing the concept of "end-point" and "equivalence point". At the end-point there is slight excess of the acid or base. At the equivalence pt, stoichiometric amounts have reacted. $\endgroup$
    – ACR
    Commented Nov 28, 2020 at 0:45
  • $\begingroup$ @M. Farooq. OK. You are right. I should have rewritten my text. But this does not change too much my message. $\endgroup$
    – Maurice
    Commented Nov 28, 2020 at 9:56

2 Answers 2

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When the $\ce{Na2CO3}$ is completely neutralized by the $\ce{HCl}$, the solution will be saturated with $\ce{CO2}$, so the pH will be lower than 7 - think of carbonated beverages. Although the $\ce{CO2}$ bubbles out, not all of it bubbles out - not even if stirred. If you had stopped the titration at pH = 7, the pH reduction from 11 would have been partly from $\ce{HCl}$ and partly from $\ce{H2CO3}$.

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  • $\begingroup$ Okay, so the low pH at the equivalence point can be attributed to just the dissolved $\ce{CO2}$, correct? $\endgroup$
    – jerrythealpaca
    Commented Nov 27, 2020 at 16:02
  • $\begingroup$ Yes. You might expect that CO2 is all expelled, but some just loves to stay in solution. $\endgroup$
    – James Gaidis
    Commented Nov 28, 2020 at 2:13
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If you have a 0.1M NaHCO3 solution with Bromophenol blue indicator (pH range 3.1-4.6) titrated with (0.1M HCl) you will get an end point (the point in a titration at which the colour of the solution shows a pronounced change) of c.80-85% of the theoretical titer or equivalence point (the point in a titration at which the amount of titrant added is chemically equivalent to the amount of sample to be titrated), ie the blue akaline solution will turn yellow more quickly than anticipated because the carbonic acid has conributed to an increase of H3O+ ions.Leaving the solution for 1-2 hours, will make it turn from yellow to green ie more alkaline as the CO2 leaves the solution under partial pressure and henrys Law.This can be shown quickly by using a water aspirator and a bung to create a vacuum in the titration vessel. Remember carbonated water is always more acidic,than distilled water which has been left to absorbed atmospheric CO2.The early end point will also be determind by how much CO2 is inside your water to start which is used for your titration solutions although this is far less than the carbonic acid contribution. The CO2 and hence the the HCO3- conc will be higher ....it is acting as a buffer and until all the excess CO2 (aq) and CO2(g) has equilibrated out of the solution into the air, the solution will remain more acidic.

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