Polverone - 12-3-2003 at 11:28
I am not sure I endeared myself to my professor today. In class we've been talking about acid/base behavior. I started wondering to myself: if
water undergoes autoionization to produce H<sub>3</sub>O<sup>+</sup> and OH<sup>-</sup>, might not other
substances undergo autoionization also? For example, HNO<sub>3</sub> dissociating to
H<sub>2</sub>NO<sub>3</sub><sup>+</sup> and NO<sub>3</sub><sup>-</sup>? Unfortunately, the
conversation soon took a major detour when I started talking about anhydrous sulfuric acid and she started saying that pure acids without water were
gases. Without saying "you are wrong" I tried to explain that many pure acids are not gases. She tried to tell me that "oleum" was
an old name for alkenes and butter! I e-mailed her the dictionary entry for oleum. I am not sure this will improve things between us.
Anyway, can anybody shed light on non-aqueous acid/base chemistry? What about sulfuric acid dissolved in acetic acid? What about nitric acid dissolved
in sulfuric acid? What about anhydrous acids undergoing their own autoionization processes?
In many textbook problems we've been asked to calculate hydronium ion concentration based on X moles of acid dissolved in Y moles of water. In
most of these problems Y >> X. But what if X > Y? I.E. what if you have a solution of 1 mole of water in 10 moles of sulfuric acid? Is
hydronium ion concentration limited by the amount of water available? Is the pH actually higher if you don't have sufficient water to make
hydronium ions from all the free acid? Do you only calculate the pH based on how much water you *do* have available?
If you pressurize H<sub>2</sub>S so that it liquifies, can you have analogous acid/base chemistry taking place in it like you would in
water?
BASF - 12-3-2003 at 12:11
A lot of interesting questions at once....
Do anhydrous acids undergo autoionization?-Yes, to a low extent.
I think entropy is the main driving force for that...
For example, sulphuric acid is known to be fully dissociated at a conc. of only 30% where it also has the biggest grade of ionization, the biggest pH
and the biggest electric conductivity.
But the concentrated acid is nearly undissociated, and thus can be stored in cheap iron containers.
In spite of that, 100% sulfuric acid DOES dissociate by itself:
2H2SO4 >H3SO4+ + HSO4-
(Ref.: CD-Roempp chemistry lexicon, 1995; Thieme Publishing)
I think the pH of acetic acid dissolved in sulfuric acid and vice-versa is a difficult question, because both can deprotonate and solvate each
other´s protons.....if that wasn´t enough, acetic acid is, like sulfuric acid, a dehydrating substance.....i would not dare to say if it was this or
that way....
BTW, if you want to calculate the pH of more concentrated solutions, it gets complicated, because then you would have to involve the activities of the
ions.....on the other hand i´m quite sure that autoprotolysis of excess H2SO4 is not that important, especially as you have to expect calculations
involving high concentrations to be FAR more inaccurate than with ideal solutions.
ooh! another question
Polverone - 12-3-2003 at 12:44
This I could probably look up for myself, since I have the references nearby, but anyway...
Okay, if pure anhydrous sulfuric acid has little ionization, then it is a poor electrical conductor, right? Likewise I would expect pure anhydrous
nitric acid to be a poor electrical conductor. But but BUT - is anhydrous nitric acid in anhydrous sulfuric acid a better conductor, because of
nitronium ion formation? Or anhydrous sulfuric acid in anhydrous perchloric acid?
And just to toss another question in here: when p-toluenesulfonic acid, or similar acids, dissociate, where does the H<sup>+</sup> come
from? I.E. which of its hydrogens does the acid give up? And why *that* particular hydrogen?
vulture - 12-3-2003 at 13:07
If you mix anhydrous nitric acid and sulfuric acid you get nitronium, oxonium and hydrogensulfate ions. If you've got ions in a solution,
you've got a conductive solution. 
I assume p-toluenesulfonic acid has an SO<sub>3</sub>H group?
If so, the H that is on the C-S-O-H bond will go into solution, simply because it's already partially positively charged by the electronpulling
effect of the oxygen.
ah, thanks
Polverone - 12-3-2003 at 13:16
I would've assumed it was the hydrogen bound to the SO3 group... if I'd noticed there was a hydrogen there in the first
place! 
Do you know of electrical conductivity studies/information that would reveal the extent of autoionization of various pure acids?
vulture - 13-3-2003 at 01:14
I haven't got any specific information on this subject but I suddenly realised something.
In pure water, one has a considerable amount of hydrogen bonding and no conduction.
Now, looking at most strong oxygen acids, which all have an OH group with the acidic function, there should be considerable hydrogen bonding in pure
form.
This would greatly restrict autoionization.