Acids and Bases

[Bond03]

I have two questions regarding acids and bases:

1. why is sulfuric acid not considered amphoteric ?

2. how come electron-withdrawing substituents increase acidity and electron-donating groups decrease acidity?? what are examples of electron-withdrawing substituents and electron-donating groups?

thank you in advance

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[tncekm]

2. how come electron-withdrawing substituents increase acidity and electron-donating groups decrease acidity?? what are examples of electron-withdrawing substituents and electron-donating groups?

thank you in advance

Lets look at triflouroacetic acid. The way I think of this is that the three Fl will pull electrons away from the electron loving oxygen that is attached to the acidic hydrogen. When its electrons get stolen by the greedy little Flourines, it will pull on the hydrogens electrons even more forcefully eventually stripping H of its electrons.

More correctly, there are two major ways to stabilize negative charges which occur as a result of the loss of an acidic proton. 1) Resonance delocalization of electrons, which is seen in carboxylic acids, and through delocaliztion of electrons by induction. Adding electronegative groups will allow for stabilization of the conjugate base by induction.

 

[bluemonkey]

Sulfuric acid, H2SO4 is polyprotic, not amphoteric. The conjugate base of sulfuric acid, bisulfate (HSO4-), however, is amphoteric. Remember that an amphoteric species can act as either an acid or as a base. H2SO4 can only donate protons, but bisulfate can either can a proton to reform sulfuric acid, thus acting as a base, or it can donate another proton to form sulfate, thus acting as an acid. Keep in mind that the pKa for this second proton is much much higher than the pKa for the first proton, thus reflecting the lower acidity. (-3 for H2SO4 vs. 2 for bisulfate).

As for the second part of your question, consider what is left over when a proton leaves an acid as illustrated by acetic acid, CH3COOH. The conjugate base of acetic acid, the acetate ion has a negative charge which is a base. The more localized the charge is on a base, the stronger the base. The reason why acetic acid is much more acid than a similar alcohol, ethanol, is because it can delocalize the negative charge via resonance. Anything that can help to further withdraw this negative charge will decrease the base strength, thus increasing the acid strength. Good examples of EWG's are strongly electronegative groups like Cl, F, etc. Keep in mind that as conjugate base strength decreases, conjugate acid strength increases via KaxKb=Kw. NO2 is another good example of an electron withdrawing group, whereas NR2 is a strongly electron donating group.

I hope this helps...

 

[iA-MD2013]

Sulfuric acid, H2SO4 is polyprotic, not amphoteric. The conjugate base of sulfuric acid, bisulfate (HSO4-), however, is amphoteric. Remember that an amphoteric species can act as either an acid or as a base. H2SO4 can only donate protons, but bisulfate can either can a proton to reform sulfuric acid, thus acting as a base, or it can donate another proton to form sulfate, thus acting as an acid.

Well...Can you really consider bisulfate to be amphoteric? Since sulfuric acid is such a strong acid, bisulfate will be a really, really weak base. We don't consider Cl- to be a base...so I don't think bisulfate has any basic activity and is simply a weak acid (not amphoteric).

 

[bluemonkey]

Actually, Cl- is considered to be a base, albeit very weak.

To the OP, I suppose H2SO4 could be considered amphoteric as it is capable of protonating itself to give H3SO4+ + HSO4−

 

[bluemonkey]

Well...Can you really consider bisulfate to be amphoteric? Since sulfuric acid is such a strong acid, bisulfate will be a really, really weak base. We don't consider Cl- to be a base...so I don't think bisulfate has any basic activity and is simply a weak acid (not amphoteric).

You can even take this one step further. Any polyprotic acid, H2A, has a conjugate base, HA-, that is amphoteric...

 

[iA-MD2013]

You can even take this one step further. Any polyprotic acid, H2A, has a conjugate base, HA-, that is amphoteric...

Well, I guess you're right in that it's amphoteric...but the definition of amphoteric is so broad, that may compounds can be considered amphoteric. For example, any amine can be protonated to make ammonium but can also be deprotonated to make an amide (NH2-). But how acidic is an amine? Can you really say that an amine is ampoteric? I tend not to chose answers that are too extreme unless it's the best choice.
I once encountered a problem that involved choosing a good buffer system. One involved H2SO4/HSO4-...Seeing as bisulfate is a really crappy base, this isn't the best buffer. I would have chosen that...but another choice was H2SO3/HS03- a weak acid/weak base system, the correct answer.

 

[bluemonkey]

Well, I guess you're right in that it's amphoteric...but the definition of amphoteric is so broad, that may compounds can be considered amphoteric. For example, any amine can be protonated to make ammonium but can also be deprotonated to make an amide (NH2-). But how acidic is an amine? Can you really say that an amine is ampoteric? I tend not to chose answers that are too extreme unless it's the best choice.
I once encountered a problem that involved choosing a good buffer system. One involved H2SO4/HSO4-...Seeing as bisulfate is a really crappy base, this isn't the best buffer. I would have chosen that...but another choice was H2SO3/HS03- a weak acid/weak base system, the correct answer.

You are right that H2S04/HSO4- is a poor buffer system. An amine is definitely amphoteric. It is capable of becoming protonated or being deprotonated. No, an amines are generally not considered very acidic, but as you mentioned, ammonia can be deprotonated into the amide ion, which does indeed mean that it can behave as an acid, hence the amphoteric classification. Granted, ammonia would not be a correct answer choice for the best acid, but it can behave as an acid.