can amides become protonated at low pH?

[JKMBC]

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I've been reviewing some recent practice questions and realized that I'm not quite sure what to do with functional groups other than carboxylic acids or amines when the pH varies. If a compound has a functional group such as an amide, do we assume that this does not act as a base and become protonated at low pH values? Is it only carboxylic acids, amine groups and phosphates that become protonated and de-protonated in compounds?


Hope this doesn't sound too dumb...

thanks!

 

[sazerac]

Amides have a pKa of -0.5, which is pretty damn acidic, and not biologically relevant.

 

[JKMBC]

So if the pKas are outside the boundaries of biological pH, they won't be affected right?

 

[popopopop]

Like sazerac said, unless the pka gets to be that low, don't worry about it being a base. Per wiki:

"Compared to amines, amides are very weak bases. While the conjugate acid of an amine has a pKa of about 9.5, the conjugate acid of an amide has a pKa around −0.5. Therefore, amides don't have as clearly noticeable acid-base properties in water. This relative lack of basicity is explained by the electron-withdrawing nature of the carbonyl group where the lone pair of electrons on the nitrogen is delocalized by resonance. On the other hand, amides are much stronger bases than carboxylic acids, esters, aldehydes, and ketones (conjugate acid pKa between −6 and −10)."

So unless it's in a solution with the above, don't worry about it acting like a base, but know it's place in reactivity.

 

[theonlytycrane]

I've been reviewing some recent practice questions and realized that I'm not quite sure what to do with functional groups other than carboxylic acids or amines when the pH varies. If a compound has a functional group such as an amide, do we assume that this does not act as a base and become protonated at low pH values? Is it only carboxylic acids, amine groups and phosphates that become protonated and de-protonated in compounds?


Hope this doesn't sound too dumb...

thanks!

The henderson-hasselbalch equation is useful in determining how a functional group (or any molecule) will act if you know the pKa of the molecule and the pH of the solution. Remember, a higher Ka or lower pKa means that a molecule is more acidic.

For example, a protonated amine group has a pKa around 9-11 (not very acidic). In a solution with pH = 3 and pKa = 9, plugging into the equation above shows that the majority of the amine group will be in acid form. If we add base to the solution until the pH > pKa, then the majority of the amine group will deprotonate into base form.