3
$\begingroup$

How was the pKa value (-1.7) for the oxonium ion determined? Why does it correspond to the minimum pH value accessible in water? What reaction does it represent?

$\endgroup$
5
$\begingroup$

There is a nice little article in the Journal of Chemical Eduction Negative pH Does Exist pointing out examples of negative pHs in water, for example pH = -3.6 in mine water in California. So -1.7 is not the lowest possible pH in water. Keep in mind that pH is defined in terms of hydrogen ion activity rather than concentration. This is emphasized in Paradoxes: Demonstrating That It Is Not True That pH ≡ -log[H+]. At high acid concentation, hydrogen ion activity greatly deviates from hydrogen ion concentration. For example, in 16 molar HCl, hydrogen ion activity is 678!

And see Negative pH and Extremely Acidic Mine Waters from Iron Mountain, California (no paywall), which describes pH values even below -4.

In Hydronium ion activity in strongly acidic media. Remarkable agreement between independent estimates Am. Chem. Soc. vol. 95, pp 3055–3057, 70% sulfuric acid is found to have a pH of -9.8 (negative 9.8)!

However, ignoring activity and the true definition of pH, considering that pure water has a concentration of 55M, if all the water molecules were instead hydronium ions, -log(55) = -1.7. Then in articles like Who Knows the Ka Values of Water and the Hydronium Ion? this value is attributed to the pKa of hydronium, with the supposed justification that: $K_a = \frac{[\ce{H3O+}][\ce{H2O}]}{[\ce{H3O+}]} = [\ce{H2O}]$

The article New point of view on the meaning and on the values of Ka(H3O+, H2O) and Kb(H2O, OH-) pairs in water is a serious consideration of the true meaning of the pKa of hydronium, and points out that the -1.7 value for the pKa of hydronium has no justification.

In conclusion, pH can be below -1.7 and the pK of hydronium isn't really -1.7.

$\endgroup$
  • $\begingroup$ How is -3.6 possible? Please elaborate. $\endgroup$ – Marko Jan 28 '15 at 14:01
  • 1
    $\begingroup$ @Marko, OK, I added to the answer, basically because hydrogen ion activity greatly deviates from hydrogen ion concentration in concentrated acids and pH is defined as -log of hydrogen ion acitivity, not concentration. $\endgroup$ – DavePhD Jan 28 '15 at 15:00
  • $\begingroup$ Sadly, can't access the: Paradoxes: Demonstrating That It Is Not True That pH ≡ -log[H+]. $\endgroup$ – Marko Jan 28 '15 at 16:16
  • $\begingroup$ @Marko As a substitute see community.asdlib.org/imageandvideoexchangeforum/2013/07/31/… which illustrates one of the paradoxes: that [H+] can be cut in half, yet hydrogen ion activity increases is a given example $\endgroup$ – DavePhD Jan 28 '15 at 16:34
  • 1
    $\begingroup$ @Marko OK, see digitalcommons.unl.edu/cgi/… should be no paywall, has pH -4 and below, I'll add it to answer $\endgroup$ – DavePhD Jan 28 '15 at 17:49
1
$\begingroup$

How was the pKa value (-1.7) for the oxonium ion determined?

Why does it correspond to the minimum pH value accessible in water water?

Note that an oxonium ion is simply an ion in which an oxygen carries a positive formal charge. Many oxonium ions exist. One of which is hydronium ion.

The hydronium ion corresponds to the strongest acid able to exist in water for significant periods of time because a stronger acid will be immediately deprotonated by some base in the system. In other words, hydronium ion can exist in equilibrium with water.

$\ce{H3O+ + H2O <=> H3O+ + H2O}$

The reaction of a stronger acid with water will however favor the products.

$\ce{HCl + H2O -> H3O+ + Cl-}$

$\endgroup$

Your Answer

By clicking “Post Your Answer”, you agree to our terms of service, privacy policy and cookie policy

Not the answer you're looking for? Browse other questions tagged or ask your own question.