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How concentrated can an acid be without it being supersaturated? Is there a certain limit to how concentrated an acid or base can be?

When I mean concentration, I mean molarity; so how concentrated can an acid or base be?


NOTE: Can be a strong or weak acid/base. Because we are just looking for the amount of solvent needed to dissolve to start the dissociation process.


EDIT: No autoprotolysis please. Acids that need a solvent to dissociate.


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    $\begingroup$ are you asking for the strongest acid in water? this is $\ce{H3O+}$. see the solvent levelling effect en.wikipedia.org/wiki/Leveling_effect $\endgroup$ – bon Mar 26 '15 at 22:06
  • $\begingroup$ In terms of theory, "how acidic can an acid be?". It depends on how much of hydrogen ions are dissociate in the water. The more the hydrogen ions dissociated in the water, the stronger the acid is which is very acidic. Examples of strong acids are Phosphorus acid, Sulphric acid, Hydrochloric acid and Hydrofluoric acid. Examples of weak acids are ethanoic acid, methanoic acid etc. $\endgroup$ – Chee King Mar 27 '15 at 2:27
  • $\begingroup$ I'm afraid you didn't really answer the question but repeated @asker12 misconception - undissociated acid is still acid and in many cases stronger than hydronium. $\endgroup$ – Mithoron Mar 28 '15 at 1:56
  • $\begingroup$ The problem is that leveling effect is here a source of misunderstanding - is important in diluted not highly concentrated solutions. $\endgroup$ – Mithoron Mar 28 '15 at 2:02
  • $\begingroup$ chemistry.stackexchange.com/questions/15421/… $\endgroup$ – Mithoron Mar 29 '15 at 13:04
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Are you suggesting that if an acid molecule is surrounded by a sufficiently small number of water molecules, then it might display lower than 100% ionization even if it is a "strong acid" ($pK_a<-1.76$)? That is true; some acids we commonly consider to be strong, such as nitric, hydrochloric or sulphuric acid, will not completely ionize in sufficiently concentrated aqueous solutions, even if there is more than one molecule of water per molecule of acid.

So is there an acid capable of protonating water quantitatively even when there is only one molecule of water per molecule of acid? Yes, in fact there are several. Interesting proof of this is that sufficiently strong acids will form solid crystalline salts containing the hydronium cation when mixed with water in a 1:1 molar ratio. The most striking examples are probably the hydronium carborane salts, such as $\ce{(H3O)^+(HCB11Cl11)^{-}}$. I recommend taking a look at this fantastic article from the Reed group for details on the wonderful properties of carborane superacids, with a section describing its several salts when reacted with water in different proportions and environments.

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I am talking about dissolving in water since an acid needs some sort of solvent to dissociate.

No, it doesn't. You neglect autoprotolysis where the water-free acid itself is the solvent.

\begin{align} \ce{2H2SO4 &<=> H3SO4+ + HSO4-}\\ \ce{3HF &<=> H2F+ + HF2-} \end{align}

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  • $\begingroup$ But where does the extra H+ come from? $\endgroup$ – Asker123 Mar 26 '15 at 19:41
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    $\begingroup$ which extra H+ ? $\endgroup$ – bon Mar 26 '15 at 22:06
  • $\begingroup$ Even without autoprotolysis lack of solvent is not a problem. $\endgroup$ – Mithoron Mar 28 '15 at 2:02

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