Why is the Bronsted-Lowry definition of acids and bases more encompassing than the Arrhenius definition?

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 . With Arrhenius bases, it is additionally specific in that a ##"OH"^(-)## must be donated to solution... while protons aren't really considered.

One example of a Brønsted-Lowry base that is NOT an Arrhenius base is sodium ethoxide (##"NaOCH"_2"CH"_3##) dissolved in ethanol (##"CH"_3"CH"_2"OH"##).

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 ##"OH"^(-)## to water; it can't, because we aren't even using water as the !

Thus, sodium ethoxide in ethanol is not an Arrhenius base; though, it IS a Lewis base since its oxygen donates two to get its proton.