Sulfuric acid as a Lewis acid

[549]

Hi,

I was doing a practice test, and one of the questions was:

Sulfuric acid can be described as all of the following EXCEPT:
a strong electrolyte
a Lewis acid
a diprotic acid
an amphoteric compound

The answer was amphoteric compound, which I got. I struggled with this one though because I did not see how it could act as a Lewis acid either.

I know that Lewis acids are electron acceptors, but the only place I thought additional electrons would go is to the sulfur, which doesn't make sense to me because if you have sulfuric acid in water, the water's not going to attack the sulfur- it's going to rip off the proton!

Also, I figured you would classify it as a BL acid instead, and I know that not all BL acids are necessarily Lewis acids.

I'd appreciate an explanation. I know it's probably not that important, but it was driving me crazy.

Thanks.

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[ratatat]

Hi,

I was doing a practice test, and one of the questions was:

Sulfuric acid can be described as all of the following EXCEPT:
a strong electrolyte
a Lewis acid
a diprotic acid
an amphoteric compound

The answer was amphoteric compound, which I got. I struggled with this one though because I did not see how it could act as a Lewis acid either.

I know that Lewis acids are electron acceptors, but the only place I thought additional electrons would go is to the sulfur, which doesn't make sense to me because if you have sulfuric acid in water, the water's not going to attack the sulfur- it's going to rip off the proton!

Also, I figured you would classify it as a BL acid instead, and I know that not all BL acids are necessarily Lewis acids.

I'd appreciate an explanation. I know it's probably not that important, but it was driving me crazy.

Thanks.

According to EK, a "Lewis acid definiton includes all of Bronsted-Lowry acids and bases and more". I dunno how that works though.

However, I heard Lewis acids are pretty much electrophiles, while Lewis bases are nucleophiles. So I guess you can think because the oxygen atoms are electronegative, they pull negative charge away from the sulfur atom. Thus, the sulfur atom can act as an electrophile???

That's my stab at an explanation. Good luck.

 

[kentavr]

I would say that first H2SO4 deprotonates. H2SO4 --> H+ + HSO4-
And then H+ is a perfect electron acceptor. (This is the definition of lewis acid)

 

[Baylor2012]

Hi,

I was doing a practice test, and one of the questions was:

Sulfuric acid can be described as all of the following EXCEPT:
a strong electrolyte
a Lewis acid
a diprotic acid
an amphoteric compound

The answer was amphoteric compound, which I got. I struggled with this one though because I did not see how it could act as a Lewis acid either.

I know that Lewis acids are electron acceptors, but the only place I thought additional electrons would go is to the sulfur, which doesn't make sense to me because if you have sulfuric acid in water, the water's not going to attack the sulfur- it's going to rip off the proton!

Also, I figured you would classify it as a BL acid instead, and I know that not all BL acids are necessarily Lewis acids.

I'd appreciate an explanation. I know it's probably not that important, but it was driving me crazy.

Thanks.

It looks a bit confusing but here is a simple explaination. Now I over analyzed it at first but after some brief thought here is my take on it. We already know that its diprotic b/c it dissociates two protons. We know its a strong electrolyte because of the fact that its a strong acid, they always dissociate. As you and others before me have stated, a lewis acid will accept a pair of electrons. But at this point it got me also. The only way I could get through would be to eliminate answer choice D b/c sulfuric acid is an acid not a base. It would only be amphiprotic/amphoteric if they'd given us say HSO4-....I hope that helps.....I'm currently looking into that lewis acid thing....because its on my nerves now also....PM me and ill get to you when i find out...

 

[549]

Okay, thanks! The above poster mentioned that the H+ would act as the electron acceptor (Lewis acid) in this case, which is the most reasonable way I've been able to look at it.

 

[Baylor2012]

could we also call hso4- as a lewis base in this case because it can donate an electron pair?

 

[ezsanche]

The answer is D. H2So4 is a lewis acid because it is a bronsted acid. ALL bronsted acids are lewis acids (but not the other way around).

H2SO4 is not amphoteric because it can not act as a base. If it were able to act as a base then you should get H3SO4 (+) as product. But I dont think that is possible so D is my final answer.

 

[thebillsfan]

The answer is D. H2So4 is a lewis acid because it is a bronsted acid. ALL bronsted acids are lewis acids (but not the other way around).

H2SO4 is not amphoteric because it can not act as a base. If it were able to act as a base then you should get H3SO4 (+) as product. But I dont think that is possible so D is my final answer.

i dont buy the whole H+ is a lewis acid deal. it is, but that's not related to the question. because we're talking about undissociated H2SO4, not a dissociated proton. Now, when sulfuric acid releases a proton, it fully accepts the pair of electrons with which it had been holding onto the proton. thus, it accepts an electron pair. Thus, lewis acid.

 

[549]

could we also call hso4- as a lewis base in this case because it can donate an electron pair?

Since HSO4- is amphoteric, it can act as a Lewis base as well as an acid.

I actually found a website that answered the original questions I had pretty well. It uses HCl as an example and states: "The whole HCl molecule is acting as a Lewis acid. It is accepting a pair of electrons from the ammonia, and in the process it breaks up. Lewis acids don't necessarily have to have an existing empty orbital."

So in the case of H2SO4, the hydrogen accepts electrons from a donating nucleophile, and in that process gets separated from the rest of he compound. This argument puts to rest the idea that perhaps its electron-accepting capabilities come from the sulfur's partial positive resulting from the many electronegative oxygens bonded to it.

Here's the rest of the site: http://www.chemguide.co.uk/physical/acidbaseeqia/theories.html

Probably not the most efficient investment of time at this point, but oh well.

 

[Baylor2012]

could we also call hso4- as a lewis base in this case because it can donate an electron pair?

Since HSO4- is amphoteric, it can act as a Lewis base as well as an acid.

I actually found a website that answered the original questions I had pretty well. It uses HCl as an example and states: "The whole HCl molecule is acting as a Lewis acid. It is accepting a pair of electrons from the ammonia, and in the process it breaks up. Lewis acids don't necessarily have to have an existing empty orbital."

So in the case of H2SO4, the hydrogen accepts electrons from a donating nucleophile, and in that process gets separated from the rest of he compound. This argument puts to rest the idea that perhaps its electron-accepting capabilities come from the sulfur's partial positive resulting from the many electronegative oxygens bonded to it.

Here's the rest of the site: http://www.chemguide.co.uk/physical/acidbaseeqia/theories.html

Probably not the most efficient investment of time at this point, but oh well.

Wow, we stretched this question a long way. But I think now we can all rest assured that we now know the logic behind it. It hit me so hard while I was driving today, almost like an epiphany!

 

[boaz]

http://forums.studentdoctor.net/showthread.php?t=655967

 

[PRATFO]

In order for it to "donate a proton" something has to donate electrons to the H

 

[wait4me]

I read somewhere that lewis acids/bases are the broadest category so all lewis acids/bases are also bronstead and arrhenius acids and bases.

Most specific--------->broadest
Arrhenius-->Bronsted--->lewis

 

[arsenal14]

So this is how I picture it.

A Lewis Base donates electron pairs, but what does this actually mean? It is not that the base is giving away its electron pairs, but rather it is sharing them with (more often than not) the Hydrogen ion which was de-protonated by the acid in the reaction. Lewis Bases have lone pairs, e.g., NH3, OH-, PH3. These lone pairs form bonds with H+ (usually).


A Lewis Acid accepts electron pairs. So what we know from bronsted acids is that B. acids donate a H+ proton. When the acid donates this hydrogen, it 'steals' the electron that belonged to the hydrogen. Lewis Acids are usually molecules which do not have a full octet. So H2SO4 when it is deprotonated steals the electrons from the ejected hydrogen ion. Another example is BF3 and AlCl3. These don't neccessarily have H+ to deprotonate, so where do they ''steal'' electrons from? If you draw their lewis structures you see that they are in need of 1 electron pair in order to fill the octet, and thus become stablilized. Lets look at an example showing a lewis acid in action. Take the first step of the Friedel Crafts alkylation reaction:

R-Cl + AlCl3 --> R+ + AlCl4-
This lewis acid catalyzed reaction allows the alkyl group to be subject to nucleophilic attack by the electron rich double bond of an arene.

I hope I didn't confuse you. All in all my basic mindset is that Lewis Acids ''steal'' electrons and Lewis Bases "share" electrons with another molecule by bonding with them. Hope this helps! If there are any discrepancies, please feel free to correct me!