What factors determine the acidity of hydrogens in vitamin C?

[ngu9997]

So in my organic chem class we learned how to choose the most acidic hydrogens on a molecule based on the stability of the conjugate base of that molecule. And then we used various factors to determine the stability of the conjugate base such as resonance (how resonance may help the negative charge from deprotonation become spread over a larger area).

For vitamin C there's three atoms that can be deprotonated to become the conjugate base. And in these 3 situations, resonance is all present. Two of the situations are an OH becoming deprotonated and one situation is a carbon becoming deprotonated. Typically CH bonds are never really seen as acidic hydrogens. Could someone explain why this is so - and in terms of the logic my class has been using would be especially appreciated (in terms of the stability of the conjugate base).

 

[TeethWhitener]

It's unclear what you're referring to. Vitamin C has 4 hydroxyl groups, only two of which can be deprotonated in water to any significant extent (pKa's of ~4 and ~11). These anions are stabilized by resonance, and some of the resonance structures you can draw contain carbanions. Is this what you're referring to? None of the carbons will be deprotonated (at least not without a much stronger base--and certainly not in water).

 

[ngu9997]

Yes - the two hydroxyl groups on the same side as the alkene in the ring are the ones that can be deprotonated creating several resonance structures for both situations. I was wondering what the reason for the hydrogens on the carbons and alkanes in general being so hard to deprotonate, specifically in terms of the logic that my class is using.

 

[TeethWhitener]

Are you asking why the conjugate base of a carbon acid is not stable?

Sometimes protons attached to carbons can be acidic, usually when the extra charge can be delocalized over a large area. One example of this is the enhanced acidity of triphenylmethane. Deprotonating at the methine moiety gives a very stable anion where the charge is delocalized across three phenyl groups.

Another important example of a carbon acid is a terminal alkyne, with pKa ~ 25 or so (most carbon acids are not strong enough to be deprotonated in water). The conjugate base here, an acetylide, is stabilized by the fact that the extra electron is in an sp orbital, which, due to its 50% s character, is less shielded from the nuclear charge than sp² and sp³ orbitals.

One final important example of carbon acid is an enolate. Ketones can be deprotonated at the carbon alpha to the carbonyl carbon. It is stabilized by resonance with the anionic enol tautomer (an enolate).

 

[TeethWhitener]

Here's a reference with a bunch of Bordwell pKa measurements in DMSO:
https://www.chem.wisc.edu/areas/reich/pkatable/index.htm
You can browse around or search by functional group to get a feel for which groups stabilize and which groups destabilize conjugate bases.