Which is more acidic?

[brood910]

---

Members don't see this ad.

So, I know that electron withdrawing groups generally make compounds more acidic. But what if the central atom is partially positive, not negative in a way that withdrawing groups DESTABILIZE it instead of STABILIZING it? I attached a picture to help you understand what I am talking about.

In this case, is the right one more acidic as it is more stable?
The central C of the left one loses even more electrons because of inductive effect by F, so it becomes even more positive and becomes less stable..

 

[WaxEarplugsFTW]

I think all that matters is the electron density around the hydrogen. The positive center is electrophilic so it tugs at the electrons near the hydrogen. Since electrons is what ties two atoms together, the "rope" is now weaker on the hydrogen side. With the added fluorine, now the positive is even more positive. Electron density around the hydrogens is even less, making the bond really really weak.
Since acidity is basically a measure of how strong the bond between hydrogen and another atom is, the left molecule would be more acidic since the bond is so weak.

I think that's how it is.
I thought stabilizing essentially means "making the charges more neutral". In that case, the right one is more stable.

Correct me if I'm wrong.

oh wait after reading what you wrote, I basically repeated what you said.

 

[brood910]

I guess they will not ask us to tell which is more acidic between these as we just have to know how to rank acidity based on the stability of conjugate bases, which always have - sign.
So, we are good as we know how to rank stability.

 

[Czarcasm]

I guess they will not ask us to tell which is more acidic between these as we just have to know how to rank acidity based on the stability of conjugate bases, which always have - sign.
So, we are good as we know how to rank stability.

They can actually ask you to rank acids in their acid form. There's two ways you can approach those types of questions. You can either look at the stability of the conjugate base (this is what I prefer), or you can look at the strength of the bond -- in its acid form -- where the hydrogen resides and distinguish things that way (less intuitive).

For the two examples you gave above, if I were to rank them both, I'd first consider their conjugate bases. Because losing a a proton results in a negative charge, the question then becomes which of the two conjugate bases are more able to stabilize this charge. The more stable conjugate base is the more acidic acid. In the example above, the group with the three fluoro groups attached is considerably more acidic because the negative charge is more withdrawn through the molecules sigma bonds towards more electronegative atoms. This is a stabilizing effect, which explains why that compound will be more acidic.

If on the other hand those fluoro groups were exchanged with three methyl groups, this would destabilize the conjugate base (making it a weaker acid), because methyl groups donate electron density via hyperconjugation.

 

[brood910]

It is easy to rank them when they are in their acid forms. As I said above, you just have to compare their conj bases.

How can you remove a proton when they are tri with a positive charge?
Even if what you said can happen, they will have a neutral charge, and there will be no inductive effect by withdrawing the electrons.

 

[Czarcasm]

It is easy to rank them when they are in their acid forms. As I said above, you just have to compare their conj bases.

How can you remove a proton when they are tri with a positive charge?
Even if what you said can happen, they will have a neutral charge, and there will be no inductive effect by withdrawing the electrons.

Sorry, but I'm not really following what you mean here.

 

[Schenker]

Let's back up all the way to the first post.

The compounds you've drawn are carbocations. Perhaps by the "+" you meant partial positive charges, but the carbocations as represented in that picture have nothing to do with the acidity of their parent compounds. Here, the relevant ionic species would be the carbanion:

The two compounds on top are the parent compounds (acid), which are neutral. The two compounds on bottom are their conjugate bases, which are negatively charged. The top-left compound is more acidic than the top-right compound due to the inductively electron-withdrawing effects of CF3 that stabilize the conjugate base.

 

[brood910]

I know how to rank acidity when they are in their normal form as all I need to do is to compare their conj bases.
So you are saying that we cannot compare carbocations for their acidity?

 

[Czarcasm]

-- edit

 

[Czarcasm]

I somehow disregarded that those examples were carbocations (I was comparing neutral species).

Regardless, a similar trend follows. In those examples above, the only hydrogen atoms available for abstraction (deprotonation) is on the neighboring alpha carbon. Both of these hyrogens reside next to an electron deficient carbon. But which carbon has a greater overall charge? To exaggerate things a bit, image a giant plus sign on the carbon with 3 fluoro groups attached. It is this carbon that wants electrons more badly (being pulled inductively by the fluoro groups acts to destabalize this charge). This region of positive charge will pull on its neighboring atoms to a stronger degree (it desperately wants negative charge), and as a result, those protons will be relatively electropositive. This alone can help you determine which species is more acidic -- but if it helps you can also consider the stability of the conjugate bases in each compound.

Consider what happens when you deprotonate -- as the proton is abstracted, two electrons are left behind. Rather than having one carbon be negatively charged and the neighboring atom positively charged, it's likely the electron pair would simply form a double bond to maintain a neutral species. In either case, the more unstable species (more reactive species) wants to form that double bond to a greater extent and for that reason, its alpha hydrogens are more likely to be deprotonated (thus making it more acidic).

Notice that in either case, whether the compound has a positive charge or not, the same trend follows.
Consider what happens if you had a carbanion instead.

I hope this helps.