- Acid-base definitions
- Organic acid-base mechanisms
- Ka and acid strength
- Ka and pKa review
- Using a pKa table
- Using pKa values to predict the position of equilibrium
- Stabilization of a conjugate base: electronegativity
- Acid strength, anion size, and bond energy
- Stabilization of a conjugate base: resonance
- Stabilization of a conjugate base: induction
- Stabilization of a conjugate base: hybridization
- Stabilization of a conjugate base: solvation
Stabilization of a conjugate base: induction
How a conjugate base is stabilized by the inductive effect.
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- Is there a good qualitative method (i.e. method other than pKa memorization) to ranking acidity of protons of acids that are stabilized by resonance compared to those that are stabilized by induction?(12 votes)
- Resonance stabilization has much more effect in stabilizing conjugate bases.
In general (Resonance effect)>(Hyperconjugation effect)>(Inductive effect)(16 votes)
- aren't inductive effect and resonance basically almost the same?
even resonance in conjugate bases(above examples) can be explained by movement of electrons of oxygen because of electronegativity. or not?(1 vote)
- inductive effect is when electrons are shifted towards more electronegative atom but stay in the same place and resonance is when they are literally delocalized over more than one atom, making new bonds or electron pairs.(14 votes)
- but having cl at a farther distance can also delocalise the entire negative charge over the anion, how is it then the fourth case is not more acidic(3 votes)
- The inductive effect is transmitted through the σ bonds.
It decreases rapidly as the number of bonds increases.
After about four bonds, the effect is close to zero.(7 votes)
- Can you help us how to visualize the term "electron density" that you keep referring in the video. I have hard time to see the importance of electron density and how it participates during these reactions.(2 votes)
- Try watching this Crash Course video on orbitals:
- Can anyone explain why more branched alkyl groups exhibit higher +I effect i dont understand it and it confuses me while watching many videos(2 votes)
- when charge is dispersed and it increases stability but more electronegative atom (oxygen) getting negative charge also increases stability.then why negative charge displacement increases stability of the conjugate base .(2 votes)
- Which compound is more acidic and why?
4-chlorobutanoic acid or
2-bromobutanoic acid? Or
- 2-Bromobutanoic acid is the most acidic.
The electron-withdrawing effect of the halogen atoms makes all of these acids stronger than butanoic acid.
They pull electron density from the O-H bond, so the H atom is held less tightly and the acid becomes more acidic.
The Cl is the most electronegative atom, but it is furthest away from the acidic H, so it has little effect.
The I atom is closer, but it is less electronegative than Cl, so it probably has little effect as well.
The Br atom is a close as it can get to the acidic H, and it is more electronegative than Br, so it has the greatest effect.(2 votes)
- So the reason the uppermost molecule (lacking Cl) is the least acidic is because it lacks comparatively electronegative atoms that can accept the surplus electron once the acid has shed a proton?(1 vote)
- can anybody please offer me a mnemonic for remembering the functional groups displaying -I effect and +I effect ?(1 vote)
- I thought the higher the pKa the weaker the acid and strong conjugate base?(1 vote)
- [Voiceover] Induction is another way to stabilize a conjugate base. So if we start with acetic acid, this is acetic acid right here, and the acidic proton on acetic acid is this one, that proton has a pKa value of approximately 4.8, so if acetic acid donates that proton, these electrons in red here are left behind on the oxygen, which gives the oxygen a negative one formal charge. So on the right, this would be the conjugate base to acetic acid. If we compare acetic acid to our next compound, this is chloroacetic acid. Notice we now have a chlorine attached to this carbon. Now I didn't draw on lone pairs of electrons for chlorine, just to make it easier to see. So the acidic proton on chloroacetic acid is this proton, so then these electrons in red are left behind on the oxygen, giving the oxygen a negative one formal charge. If we look at the pKa value, the approximate pKa value for this proton, it's about 2.9, so think about the difference in acidity between chloroacetic acid, and acetic acid. You're going from 4.8 to 2.9. The lower the pKa, the more acidic the compound. And that's approximately 100 times more acidic. This is two pH units, from 4.8 to 2.9 is pretty close to two pH units. So that would be 10 to the second power, or 100 times more acidic. And so chloroacetic acid is much more acidic than acetic acid. And if we look at the conjugate bases, we can understand why. So this, this conjugate base must be more stable than this conjugate base. And we could explain this in terms of induction. So if we look at the difference, we know we have chlorine here. And chlorine is an electronegative element. It's much more electronegative than carbon, so chlorine's going to withdraw some electron density this way, and if you withdraw electron density, you delocalize this negative charge. You spread out this negative charge, and that stabilizes the conjugate base. And since this conjugate base is more stable than this conjugate base, chloroacetic acid is more likely to donate its proton than acetic acid. And we can see what happens as we increase the number of chlorines. So down here is trichloroacetic acid, we have three chlorines. And the pKa has lowered even more. Because now we have all these chlorines here withdrawing some electron density, so all these electron withdrawing groups if you will, are withdrawing electron density and that's stabilizing this conjugate base, that's spreading out this negative charge. So this is the most stable conjugate base out of these three. Therefore, this is the most acidic compound out of these three. The inductive effect falls off over distance. So if we look at this acid here, this is called butanoic acid. So this is carbon one, this is carbon two, this is carbon three, and this is carbon four. So this proton has a pKa value of approximately 4.8. If we compare butanoic acid to chlorobutanoic acid, so this would be with a chlorine on carbon two, so two chlorobutanoic acids, the pKa value has dropped to 2.8. So again, that's because of the inductive effect. We have an electronegative atom withdrawing electron density, stabilizing our conjugate base, therefore lowering the pKa value for this proton. Now if we move the chlorine to the third position, this is carbon one, this is carbon two, this is carbon three, so now the chlorine's on the third position so this is three chlorobutanoic acid. The pKa value is still lower than 4.8. 4.8 was the original pKa value for this proton. Now it's 4.1, but notice, it's not as low as it was in the previous example, so with the chlorine on carbon two, the pKa value is 2.8. When the chlorine's on carbon three, the pKa value is 4.1, so still more acidic than the original butanoic acid. But you can see, the electronegative atom is further away from the negative charge, and that decreases the effect. The chlorine is further away from the negative charge than in this example. And finally, you can see the effect even more. If you move the chlorine to the fourth position, this is carbon one, two, three and four, if you move the chlorine to the fourth position, now your pKa value is almost back up to 4.8, so it's about 4.5, so this chlorine is further away from this negative charge and that decreases the inductive effect. And so this compound is still a little bit more acidic than butanoic acid, but the effect is greatly diminished.