Carboxylic acid strengths

Here's an interesting question, let's see if anyone gets it :

Why is trichloroacetic acid about ten thousand times more acidic than acetic acid?

Because fluorine is so damn electronegative, that the negative charge 'on' the oxygen is sucked through the molecule to the fluorines, with the net effect being the negative charge delocalized throughout the fluoroacetate ion, and as delocalized charges are more stable the trifluoro acetate ion is much moe stable than acetate. Greater stability of conjugate base= stronger acid.
EDIT: Cmon, that was easy

[Edited on 6-7-2006 by rogue chemist]

Edit 2: he said chloroacetic, not fluoroacetic. But same idea.

Tim

Rogue: that has an effect, but it is not the major one .

Not so easy!!

Edit: your clue is the letter "S".



[Edited on 6-7-2006 by Nick F]

Quote:

Originally posted by Nick F Rogue: that has an effect, but it is not the major one . Not so easy!! Edit: your clue is the letter "S".

Really? I've always thought that the electron withdrawing chlorines (or fluorines ) cause the proton to be easily abstracted, leading to the increased acidity of the acid.

I'm not sure what the clue is supposed to represent though, I can't think of anything that would make a compound more acidic that starts with an "S"

Yeah, I am confused about what the reason could be now, what I said above got me full marks on a question on an o-chem final question, I also looked in my text, and it said the same.

Now you got me all curious

Actually it's not either or, its both. I merely explained why electron-suckers like halogens increase acidity

Hehehe .

To be fair, it would have taken me a very long time to come up with the answer if I had not been taught it.

S - entropy!! Not the most obvious of clues, I admit...

When acetic acid is deprotonated, the negative charge is delocalised onto only two oxygen atoms, with little effect from the CH3. Thus that part of the molecule has a high charge density, and so water molecules are very strongly attracted to it. So, when acetic acid is deprotonated, lots of water molecules are suddenly bound quite strongly to the COO- end of the molecule, which leads to a large decrease in entropy.

When trichloroacetic acid (trifluoro, etc) is deprotonated, the inductive effect from the halogens causes the negative charge to be less localised. In turn, this means that the surrounding water molecules are less strongly attracted, and so the decrease in entropy is much smaller, and so deprotonation of trihaloacetic acids is more favourable.

The inductive effect does of course also mean that the proton is more labile, but that alone is not nearly enough to account for the ten thousand-fold increase in acidity.

Told you it was interesting .

Organic chemistry - you've gotta love it!

Edit: I said "water atoms" by mistake. But it's late and I'm drunk .

[Edited on 6-7-2006 by Nick F]

Heh, interesting, thanks for the little lesson, I take physical chemistry next year. Kinda looking forward to it, a bit worried though as I seem to be the weakest in it out of all the subclasses of chemistry.

(I never really liked entropy, and I absolutly despise gibbs free energy)

Once you get the hang of the delta G it's really not to hard