# I tested the effect of temperature on the pH of carbonated water but the results seem to go opposite to what was expected. Could anyone explain?

We carried out an experiment to investigate the effect of temperature on the acid carbonate equilibrium in carbonated water. We tested 5 different temperatures (room temp, 2 above and 2 below) by heating carbonated water in a closed test tube. After leaving the test tube for 10 minutes in the water bath, we opened it and quickly recorded the temperature and pH of the solution using probes.

We got the results that as temperature increases the pH decreased (so the solution got more acidic) and vice versa, however published results show that the pH should have increased for an increase in temperature. I understand this is because the increase in temperature decreases the solubility of carbon dioxide, decreasing the amount of carbonic acid and hydronium ions in solution and therefore increasing the pH.

The entire system should have shifted as per the equation:$$2CO_{2(g)}+3H_2O_{(l)}\rightleftharpoons CO_{2(aq)}+2H_3O^{+}_{(aq)}+CO^{2-} _{3(aq)}$$

I originally thought that the opposite trend in my data was due to the endothermic nature of the ionisation of carbonic acid making the system overall endothermic, but published data indicates that the system is overall exothermic.

It would be great if anyone had any ideas about how to explain the opposite shift in equilibrium that our results show.

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• In carbonated water, content of $\ce{CO3^2-}$ is negligible, as it occurs in very alkaline solutions(50 % of it at pH near 10.3), while the water is acidic. Dec 9, 2022 at 8:49
• Did you account for the temperature dependence of $K_w$? Dec 9, 2022 at 12:44