March 28, 2024, 04:55:27 AM
Forum Rules: Read This Before Posting


Topic: Lewis Acid/Base and Redox Reactions  (Read 6324 times)

0 Members and 1 Guest are viewing this topic.

Offline galpinj

  • Full Member
  • ****
  • Posts: 102
  • Mole Snacks: +2/-5
Lewis Acid/Base and Redox Reactions
« on: August 03, 2016, 12:54:57 AM »
Hey guys,

So there are plenty of posts explaining the difference between a Lewis acid/base and a redox reaction, namely that, unlike Lewis acid/bases, a redox reaction will involve a change in the oxidation number for the specific elements.

I have, however, been unable to find out the reason why we bother specifying this difference. Why don't we call all redox reactions Lewis acid/base reactions? Is there something particular about the redox reaction that demands they be singled out?

Offline orthoformate

  • Full Member
  • ****
  • Posts: 133
  • Mole Snacks: +14/-4
Re: Lewis Acid/Base and Redox Reactions
« Reply #1 on: August 03, 2016, 01:39:15 AM »
would you draw or write out an example? Lewis acids/bases accept or donate electron pairs to another molcule, but the electron donor is always still attached to the electrons it is donated, creating a complex.

Redox reactions involves electron transfer, the electron donor is not still attached to the electrons it has donated.

Offline galpinj

  • Full Member
  • ****
  • Posts: 102
  • Mole Snacks: +2/-5
Re: Lewis Acid/Base and Redox Reactions
« Reply #2 on: August 03, 2016, 10:33:14 AM »
would you draw or write out an example? Lewis acids/bases accept or donate electron pairs to another molcule, but the electron donor is always still attached to the electrons it is donated, creating a complex.

Redox reactions involves electron transfer, the electron donor is not still attached to the electrons it has donated.

What about something like HCl --> H+ + Cl-? This is definitely a lewis acid that dissociates (though not a redox reaction).

I think the vast majority of lewis acid/base reactions still involve the electron donor and electron acceptor being separated.

Also, on a related note, can someone help clarify how we define a lewis acid and a lewis base. If we were to look at the above reaction strictly from a lewis acid/base point of view, HCl is considered an acid because the Chlorine will accept a pair of electrons from the Hydrogen. In this example, we focused on the more electronegative term/the second term (Cl) when deciding whether the molecule HCl was an acid or base.

However, if we look at a different reaction, like NaOH --> Na+ + OH-, we will call NaOH a base. This certainly makes intuitive sense, but it doesn't line up with what we normally do. In NaOH, it is the second term/more electronegative term (OH) that accepts electrons (just like Cl) while Na donates electrons (much like hydrogen in the earlier example). This makes me want to call NaOH an acid, with the resulting Na+ an acid and OH- a base.

Basically, when using the lewis acid/base definition, which part of the molecule do we look at when deciding whether the molecule is an acid or a base?

P.S. I'm aware that because HCl and NaOH are strong acids and bases, respectively, their conjugates would not really function as a base/acid. Just, hypothetically, and based on the lewis acid/base definition, how do we define the molecule (before it separates!) as an acid or base?
« Last Edit: August 03, 2016, 11:21:42 AM by galpinj »

Offline orthoformate

  • Full Member
  • ****
  • Posts: 133
  • Mole Snacks: +14/-4
Re: Lewis Acid/Base and Redox Reactions
« Reply #3 on: August 03, 2016, 12:18:46 PM »
HCl + H2::equil:: H3O+ + Cl-

this is a Bronsted acid. there is proton transfer between molecules, but there is no electron transfer.

It sounds like you understand the difference between acid/base and redox. Now you are curious about the difference between lewis and bronsted acids and bases?

https://en.wikipedia.org/wiki/Lewis_acids_and_bases
https://en.wikipedia.org/wiki/Br%C3%B8nsted%E2%80%93Lowry_acid%E2%80%93base_theory

read these two articles, they will help you differentiate between the two.

Offline galpinj

  • Full Member
  • ****
  • Posts: 102
  • Mole Snacks: +2/-5
Re: Lewis Acid/Base and Redox Reactions
« Reply #4 on: August 03, 2016, 12:49:42 PM »
HCl + H2::equil:: H3O+ + Cl-

this is a Bronsted acid. there is proton transfer between molecules, but there is no electron transfer.

It sounds like you understand the difference between acid/base and redox. Now you are curious about the difference between lewis and bronsted acids and bases?

https://en.wikipedia.org/wiki/Lewis_acids_and_bases
https://en.wikipedia.org/wiki/Br%C3%B8nsted%E2%80%93Lowry_acid%E2%80%93base_theory

read these two articles, they will help you differentiate between the two.

Thank you for the prompt response. Even after looking at the articles, I still find it odd not to exclude HCl as a lewis acid. I certainly agree that it is a Bronsted-Lowry acid (donates a proton), but it also looks like a lewis acid (the more broad and inclusive definition). Isn't the Hydrogen in HCl donating a pair of electrons while Chlorine accepts a pair of electrons (thus acting like a lewis acid)?

Offline orthoformate

  • Full Member
  • ****
  • Posts: 133
  • Mole Snacks: +14/-4
Re: Lewis Acid/Base and Redox Reactions
« Reply #5 on: August 03, 2016, 01:58:46 PM »
looks like you are right: this was taken from the wikipedia page

H+ as Lewis acid
The proton (H+) [4] is one of the strongest but is also one of the most complicated Lewis acids. It is convention to ignore the fact that a proton is heavily solvated (bound to solvent). With this simplification in mind, acid-base reactions can be viewed as the formation of adducts:

H+ + NH3 → NH4+
H+ + OH− → H2O

I guess I don't know this topic very well. maybe someone else can answer this question for you.

Offline galpinj

  • Full Member
  • ****
  • Posts: 102
  • Mole Snacks: +2/-5
Re: Lewis Acid/Base and Redox Reactions
« Reply #6 on: August 03, 2016, 07:17:55 PM »
Nooo that's not what I was hoping to hear! Thank you so much nonetheless for all the help, very much appreciated!

Sponsored Links