Could a Bronsted-Lowry acid not be an Arrhenius acid? Explain.


Could someone explain the answer to me?

Certainly! To understand why a Bronsted-Lowry acid may not be an Arrhenius acid, we first need to know the definitions of these two acid concepts.

In the Arrhenius definition, an acid is a substance that donates hydrogen ions (H+) when dissolved in water. For example, HCl (hydrochloric acid) dissociates in water to produce H+ ions and Cl- ions.

On the other hand, in the Bronsted-Lowry definition, an acid is a substance that donates a proton (H+) to another substance. Protons are commonly associated with hydrogen ions. This definition is more general than the Arrhenius definition because it is not limited to the presence of water.

Now, to answer your question: Yes, it is possible for a Bronsted-Lowry acid to not be an Arrhenius acid. This is because the Arrhenius definition is limited to aqueous solutions, while the Bronsted-Lowry definition applies to any type of solvent or reaction medium.

For example, acetic acid (CH3COOH) is a Bronsted-Lowry acid because it readily donates its proton (H+) to another substance. However, acetic acid in its pure form does not readily dissociate into H+ ions when dissolved in water, so it does not meet the criteria for an Arrhenius acid.

In summary, a substance can be classified as a Bronsted-Lowry acid if it donates a proton, while an Arrhenius acid specifically refers to substances that donate hydrogen ions in aqueous solutions. Therefore, a Bronsted-Lowry acid may not meet the criteria to be considered an Arrhenius acid depending on the conditions and solvent being used.