Chemical Forums
Chemistry Forums for Students => Organic Chemistry Forum => Topic started by: Swedish Architect on December 03, 2013, 01:30:03 PM
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I presume that the oxygen (double bonded to carbon) in carboxylic acids acts to cause this. Does it pull charge away from carbon, and by the inductive effect, pull electron density from the oxygen-hydrogen bond??
If the oxygen (double bonded to carbon) was on another atom, presumably the molecule wouldn't be acidic?
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Please can you define "acidic"
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There needs to be clarification here. As archer said, what do you mean acidic? Arrhenius? Lewis? ect..?
Start with pKa and pH definitions.
A proton is almost always labile when heteroatoms are around.
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I presume that the oxygen (double bonded to carbon) in carboxylic acids acts to cause this. Does it pull charge away from carbon, and by the inductive effect, pull electron density from the oxygen-hydrogen bond??
If the oxygen (double bonded to carbon) was on another atom, presumably the molecule wouldn't be acidic?
I think you will find a similar effect, but reduced because oxygen is more electron withdrawing than nitrogen or carbon (oxygen has a larger nuclear charge). An amide RCONH2 will be more acidic than an R'NH2 and an RCOCH3 will be more acidic than an R'CH3.
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There is a notable alcohol that is pretty acidic, by the Arrhenius definition. Comparable to carboxylic acids. Do you know which alcohol I mean? It will give yo a big hint as to what makes the difference.
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I presume that the oxygen (double bonded to carbon) in carboxylic acids acts to cause this. Does it pull charge away from carbon, and by the inductive effect, pull electron density from the oxygen-hydrogen bond??
If the oxygen (double bonded to carbon) was on another atom, presumably the molecule wouldn't be acidic?
carboxylic acids are more acidic than corresponding alcohols due to resonance stabilisation of the oxygens on the carbocation when the proton is lost
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I presume that the oxygen (double bonded to carbon) in carboxylic acids acts to cause this. Does it pull charge away from carbon, and by the inductive effect, pull electron density from the oxygen-hydrogen bond??
If the oxygen (double bonded to carbon) was on another atom, presumably the molecule wouldn't be acidic?
carboxylic acids are more acidic than corresponding alcohols due to resonance stabilisation of the oxygens on the carbocation when the proton is lost
When a proton is lost from a carboxylic acid you get a carboxylate anion, RCOO-. So there is no carbocation.
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Yep, carboxylic acids are more acidic from an Arrhenius perspective, because the negative charge is delocalized between the 2 oxygen atoms.
There is a notable alcohol that is pretty acidic, by the Arrhenius definition. Comparable to carboxylic acids. Do you know which alcohol I mean? It will give yo a big hint as to what makes the difference.
I know the question wasn't directed at me, but I'll bite anyway.
First I thought about something like tert-butyl alcohol, but thinking about it some more it's probably something like trichloromethanol, if it even exists. :P
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First I thought about something like tert-butyl alcohol, but thinking about it some more it's probably something like trichloromethanol, if it even exists. :P
I'm guessing he's referring to phenol, AKA carbolic acid, which is acidic (Arrhenius definition) for the same general reason that carboxylic acids are.
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I'm guessing he's referring to phenol, AKA carbolic acid, which is acidic (Arrhenius definition) for the same general reason that carboxylic acids are.
Yeah, phenol has a pKa of 9 but I don't get why it's so acidic relative to other alcohols. Aren't the pi bonding MO's in benzene fully occupied and so the negative charge on the oxygen would delocalize to pi anti-bonding orbitals?
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Yeah, phenol has a pKa of 9 but I don't get why it's so acidic relative to other alcohols. Aren't the pi bonding MO's in benzene fully occupied and so the negative charge on the oxygen would delocalize to pi anti-bonding orbitals?
The -ve charge of phenoxide is resonance stabillised by the aromatic ring.
Look at the pKa of substituted phenols with electron withdrawing and electron donating substituents.
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The -ve charge of phenoxide is resonance stabillised by the aromatic ring.
Look at the pKa of substituted phenols with electron withdrawing and electron donating substituents.
Phenols with electron withdrawing substituents lower pKa's, because then the delocalization of negative charge would be greater by the aromatic ring?
Yes, I know the explanation usually given is the stabilization of the negative charge by the aromatic ring, so oxygen doesn't get a full negative charge. But why doesn't this disrupt the aromaticity of benzene? And what happens to the C-C bond length in the ring when phenoxide is formed?
Hope I'm making my point clearly, sorry.
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It's true the negative charge will likely go into a delocalized antibonding orbital, and thus decrease the average C-C bond order, but so what? This is still more favorable than the electron going into a localized antibonding orbital, or being stuck in one high energy atomic orbital. If you are really curious about how this impacts the C-C bond order, you can do a simple Huckel treatment probided you have a little skill at matrix math or have access to a good math program like Mathematica.
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It's true the negative charge will likely go into a delocalized antibonding orbital, and thus decrease the average C-C bond order, but so what? This is still more favorable than the electron going into a localized antibonding orbital, or being stuck in one high energy atomic orbital.
Yeah, you're right. I hadn't thought about it like that.
If you are really curious about how this impacts the C-C bond order, you can do a simple Huckel treatment probided you have a little skill at matrix math or have access to a good math program like Mathematica.
I am. Thanks for the suggestion.
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Using the definition of an acid as a "substance which donates protons (hydrogen ions) to other things", the carboxylic acids are acidic because of the hydrogen in the -COOH group.
In addition, electronegativeee substituents near the carboxyl group actt to increaseee thhe accidity..
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The -ve charge of phenoxide is resonance stabillised by the aromatic ring.
Look at the pKa of substituted phenols with electron withdrawing and electron donating substituents.
Despite the risk of being flagged a whole lot more times for my opinion on this, I shall persist.
If asked for a quiz or exam, Archer's answer is the one to give.
If wondering about chemistry, I think we should be more skeptical. I am wary of product stabilization being a good reason to explain acidity if I cannot explain it a priori. We should ask whether the HO of phenol is different than other HO groups. For example, the pKa of H2O2, HOCl, choline, CF3CH2OH, etc are all reduced without resonance. The pKa of cyanic acid is quite low, and while it could be given the resonance argument, it also fits well with an increased electron withdrawing effect of a CN group.
Although the acidity of phenol can be explained by a resonance effect (and is the preferred reason for an exam), can we find additional data to support the resonance structures, for example in the UV spectrum. If we argue the greater contributor has the charge on the oxygen, then we should concede that perhaps the resonance structures may not have that important contribution.
The alternate feature to be examined should be the magnitude of the electron withdrawing effect of an sp2 carbon (to carbon). Herein we have lots of examples. Besides the increased acidity of ethylene and the further increase in acidity of acetylene, I argue this effect can be transmitted through atoms, such as an OH or NH. If a CN group is simply an electron withdrawing element attached to an OH in cyanic acid, then its increased acidity can be similar to HOCl. I argue we must be careful to distinguish between expecting a resonance effect and whether it is really present. The acidity of the NH of pyrrole is also increased. While we may be tempted to draw a resonance structure, let me point out that pyrrole is aromatic by virtue of the non-bonded electron participating in the ring current of the aromatic system. The NH bond electron are perpendicular to the aromatic electrons.
This effect can be seen in many five-membered aromatic heterocycles. If the neighboring atom is even more electron withdrawing, such as nitrogen, the acidity can be further increased. Compounds like benzotriazole are quite acidic and it seems to me the increase in acidity can occur without invoking an a priori resonance stabilization. If the benzene ring and sp2 nitrogen atoms can be electron withdrawing, they can inductively increase the acidity.
In anticipation of the flags I am about to receive, I am not saying resonance does not play a roll. I am saying I am wary of arguing acidity should be increased due to stabilization of an anion as if the inductive effects of the groups attached to the acid can or should be ignored. I think this was the original poster's comment and one to which I would agree.