Why is the Bronsted-Lowry definition of acids and bases more encompassing than the Arrhenius definition?
1 Answer
Well, let's put it this way:
The Brønsted-Lowry definition, being less specific, is more encompassing than the Arrhenius definition.
Here's what I mean.
The Brønsted-Lowry definition says that:
- An acid donates a proton (
#"H"^(+)# ). No further qualification is involved. - A base accepts a proton (
#"H"^(+)# ). No further qualification is involved.
The Arrhenius definition says that:
- An acid donates a proton (
#"H"^(+)# ) with the qualification that it occurs upon dissociation and the proton is donated to water. - A base donates an
#"OH"^(-)# with the qualification that it occurs upon dissociation and the#"OH"^(-)# is donated to water.
As a result of the more specific nature of the Arrhenius definition, it is confined to only aqueous solutions. With Arrhenius bases, it is additionally specific in that a
One example of a Brønsted-Lowry base that is NOT an Arrhenius base is sodium ethoxide (
We should notice that it can accept a proton (by donating electrons), just like the Brønsted-Lowry base definition requires, but it does not donate an
Thus, sodium ethoxide in ethanol is not an Arrhenius base; though, it IS a Lewis base since its oxygen donates two valence electrons to get its proton.
