When I prepared a solution with 100 ml of 0.2M KCl and 0.5 ml of 0.2M HCl, the pH of the solution was 2.33. However, when I heated it up to 50 C then the pH dropped to 0.88.
Now since HCl is a strong acid it is anyway dissociated completely even at room temperature and acidity of water doesn't increase so drastically by heating it up.
Then why does the pH of the solution go down so drastically?
Now I just want to know if there is any theoretical basis for such results?
The pH of of 0.2 M HCl aq. is pH = -log10(0.2) = 0.7. Upon dilution of 0.5 ml of 0.2 M HCl aq. sol. by adding to 100 ml 0.2 M KCl aq., the HCl molarity of the resulting diluted sol. should be approx. 0.2·0.5/100.5 = 9.95·10-4 M. For this last sol. you should have approx. pH = -log10( 9.95·10-4) = 3.0, since KCl is a neutral salt, i.e., it can be obtained by the neutralization of a strong acid (HCl) with a strong base (KOH).
On the effect of the temperature on acid-base equilibria of aq. sol., you may refer to: Herbert A. Laitinen, Walter E. Harris, “Chemical Analysis ─ An Advanced Text and Reference”, 2nd ed., McGraw-Hill, 1975, pp. 47-48.
Note that the vapour pressure of HCl above the aq. KCl/HCl sol., is quite significant and increases noticiably with the temperature, hence possibly leading to noticeably evaporation of HCl from the sol..
Also note that the pH measurements considered should be temperature compensated.
Carbonation of the KCl salt or of its aq. sol. by atmospheric CO2 may possibly play a role.
Thank you Carlos for the input. I will check out the literature mentioned as well.
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