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Topic: Stability of acetic acid and its conjugate base!  (Read 21691 times)

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Offline Winga

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Stability of acetic acid and its conjugate base!
« on: July 24, 2004, 01:05:43 AM »
Acid-base reaction:

CH3COOH  +  H2O  <===>  CH3COO-  +  H3O+
___acid__________________conjugate base_______

I found 2 explanations about the shift of equilibrium.

1.
The pKa value of hydronium ion (H3O+) is -1.7, while acetic acid is 4.7.
H3O+ is more acidic than acetic acid, thus, the equilibrium is shift to left.

2.
By comparing the stability of acetic acid and acetate, because there is delocalization of electrons presented in acetate between O-C-O atoms, acetate gains extra stability, so, the equilibrium shifts to the side (right) with more stable product.

Are there any problems in these two explanations?

Thanks!

By the way, I have a question about comparing the stability of acetic acid and acetate. In acetate, because of delocalization of electrons, the -ve charge is being diminished by resonance effect. However, acetate is still got a -ve charge although it is smaller than 1.

On the other hand, acetic acid is neutral (no charge), so, is acetic acid more stable than acetate?

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Re:Stability of acetic acid and its conjugate base!
« Reply #1 on: July 26, 2004, 12:16:09 PM »
Acetic acid is more stable because it has more bonds.  I think that the take home message from your second explanation should be that the charge on an deprotonated carboxylic acid is more stable than, say, the charge on a deprotonated alcohol.  Your two explanations are somewhat conflicting, right?  There is some population of the right hand side of the equation, but the equilibrium is shifted mostly to the left by the reasoning in your first point.  The second point means that the right side of the equation is more accessible than it would be if there wasn't that stabilization of the negative charge.

I hope that makes sense.

Offline Winga

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Re:Stability of acetic acid and its conjugate base!
« Reply #2 on: July 26, 2004, 11:51:16 PM »
But, if I compare the stability of phenol and phenoxide, both have resonance effect, and phenoxide is more stable than phenol, because in phenol, energy is needed to separate unlike charges.

This time, which one is more stable, phenol and phenoxide?


By the way, why there is charges separation in phenol (I know the mechanism)? Unlike charges should be attracted together naturally, so, is resonance effect really existed in phenol?

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Re:Stability of acetic acid and its conjugate base!
« Reply #3 on: July 27, 2004, 12:20:44 PM »
Phenol is still more stable than phenoxide, again because there are more bonds.

Similar to the acetate case, the energies are brought closer together because in phenoxide the charge can be "smeared out" over the aromatic ring.  The resonance effect in phenol is smaller than in phenoxide because it requires a positive charge on the oxygen and a negative charge somewhere else in the ring.  That is what is meant by charge separation.  It takes energy to pull those charges apart since, as you stated, they would be attracted to one another.

Offline Winga

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Re:Stability of acetic acid and its conjugate base!
« Reply #4 on: July 28, 2004, 07:06:06 AM »
From this reaction:
CH3COOH  +  H2O  <=>  CH3COO-  +  H3O+

In this case, if I add much water into the system, will the equilibrium shifts to right hand side?

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Re:Stability of acetic acid and its conjugate base!
« Reply #5 on: July 28, 2004, 12:25:21 PM »
According to LeChattlier, yes it should.

You could justify it physically by imagining that there are now more water molecules to stabilize the positive and negative ions that are created by the dissociation of the acetic acid molecules.

I think the effect would be pretty small in practice though, since there is usually a large excess of water around anyway.  I haven't worked out the math on that, but that's my gut feeling.

Offline Winga

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Re:Stability of acetic acid and its conjugate base!
« Reply #6 on: July 28, 2004, 10:44:42 PM »
I would like to know that water does not involve in the acid dissociation constant, Ka, is it because water is excess?

And why someone said the concentration / density of water does not change? (but it is true that water has changed its concentration in the reaction)
« Last Edit: July 28, 2004, 11:59:29 PM by Winga »

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Re:Stability of acetic acid and its conjugate base!
« Reply #7 on: July 29, 2004, 12:20:40 AM »
If I recall correctly, water is involved in the pKa calculation.  However when you calculate the equilibrium for a typical acid dissociation the concentration of water is essentially constant.  The change is so small that is irrelevant.  So basically yes, it is because water is in such a large excess.

For your second question, the density of water is dependent on temperature.  The concentration does change, just not very much.

Offline Winga

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Re:Stability of acetic acid and its conjugate base!
« Reply #8 on: July 29, 2004, 05:12:29 AM »
Density of water is depended on temperature because of changes in packing pattern / efficiency with temperature (e.g. ice has a lower density than water, 4 Degree Celsius has the highest density).

In a mixture, water mixed with its immiscible liquid, that both have same amount, let's say 200 mL each. Because they are not miscible, two layers of liquids are formed and now, water molecules only have interaction between themselves, so the packing of water molecules is still the same as pure water.

But when water mixed with its miscible liquid and each of them has the same volume, 200mL. There is interaction between water molecules and the molecules of that liquid, let's say "liquild A". The packing of water molecules is not as same as the pure water anymore, because there are liquild A molecules existed between water molecules. Now, the density of water is equal to its mass (200g) over total volume (400mL). Therefore, the density of water becomes half of its original density. Am I correct?

(2 cases are undergone in constant temperature)  
« Last Edit: July 29, 2004, 05:15:49 AM by Winga »

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Re:Stability of acetic acid and its conjugate base!
« Reply #9 on: July 29, 2004, 12:21:27 PM »
I honestly don't know the answer for that scenario.  My instincts tell me that density isn't really the right term for that sort of system.  Wouldn't it adhere to some sort of partial pressure law?

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Re:Stability of acetic acid and its conjugate base!
« Reply #10 on: July 29, 2004, 10:51:41 PM »
I used density because someone said that the concentration and the density are related to each other by doing some calculations. (e.g. density = mass / volume)

Is this method not quite suitable to explain the (nearly) unchanged concentration of water?

Thanks a lot!

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Re:Stability of acetic acid and its conjugate base!
« Reply #11 on: July 30, 2004, 02:02:14 AM »
Yeah, I guess that relationship between density and concentration is valid, but I think it would only really be valid if the substance were pure.

I think your explanation for the unchanged concentration of water is okay, since the amount of acetic acid in solution is probably relatively small.

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Re:Stability of acetic acid and its conjugate base!
« Reply #12 on: July 31, 2004, 10:42:30 AM »
You said acetic acid is more stable because it has more bonds, you meant the sigma bond, right?

I found that benzene is more stable than cyclohexadiene, but cyclohexaidene has more bonds.

Benzene gains extra stability because there are delocalization of electrons in the ring (resonance effect). Acetate also has resoanace effect. So, can I determine the stability of acetic acid and acetate (phenol and phenoxide ion) just by the number bonds?

Is there any alternative method?

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Re:Stability of acetic acid and its conjugate base!
« Reply #13 on: July 31, 2004, 03:11:40 PM »
The resonance effect in acetate or phenoxide are quite different than that of benzene.  The biggest difference is that acetate and phenoxide are anions, which tend to be higher in energy than neutral molecules.

I'm guessing that you mean that benzene is more stable than cyclohexaTRIene.  That is because of the resonance effect, but those are both neutral molecules.  You're kind of comparing apples and oranges.  A closer comparison would be between an acetate anion where all the charge was fixed on one oxygen (no resonance) versus the actual case where the charge can be smeared out over both oxygens.  However, this has no effect on the protonated acid stability.  That O-H bond has to be worth something, right?  The resonance effect of a carboxylic acid just makes the deprotonated form lower in energy.

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Re:Stability of acetic acid and its conjugate base!
« Reply #14 on: July 31, 2004, 07:48:30 PM »
Acid-base reaction:

CH3COOH  +  H2O  <===>  CH3COO-  +  H3O+
___acid__________________conjugate base_______

I found 2 explanations about the shift of equilibrium.

1.
The pKa value of hydronium ion (H3O+) is -1.7, while acetic acid is 4.7.
H3O+ is more acidic than acetic acid, thus, the equilibrium is shift to left.

2.
By comparing the stability of acetic acid and acetate, because there is delocalization of electrons presented in acetate between O-C-O atoms, acetate gains extra stability, so, the equilibrium shifts to the side (right) with more stable product.

Are there any problems in these two explanations?

Thanks!

By the way, I have a question about comparing the stability of acetic acid and acetate. In acetate, because of delocalization of electrons, the -ve charge is being diminished by resonance effect. However, acetate is still got a -ve charge although it is smaller than 1.

On the other hand, acetic acid is neutral (no charge), so, is acetic acid more stable than acetate?

First of all, we are refering to acetic acid in aqueous solutions...acetic acid is considered a moderately strong acid (weak relative to some other acids) and one can see this from different perspectives.  

The conjugate base is stabilized by resonance, the negative charge is somewhat delocalized and thus it has decreased reactivity as a base, to react in the reverse direction.  

Also the acidic hydrogen is bonded to an oxygen group, increasing the charge separation and thus facilitating the reaction; free energy of activation.




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