Use the Bronsted-Lowry definitions to identify the two conjugate acid-base pairs in the following acid-base reaction:
H20 + H20 <-> H30^+ + OH^-
Let's see what you think on this after the previous post. Just remember, the acid is the one that HAS the H and the base is the one that TAKES the H.
But they all have an H in them? How would I know then, which one is the acid and which one is the base?
I agree this is not that clear cut; however, the H3O^+ is a clue. That MUST have come from H2O; therefore, one of the H2O molecules gave a H to the other one to form H3O^+ and that leaves a OH^- ion.
I think the point of this question is to show you that H2O can be an acid while H2O is being a base.
To identify the conjugate acid-base pairs in the given acid-base reaction, we can apply the Bronsted-Lowry definitions. According to these definitions, an acid is a species that donates a proton (H+) and a base is a species that accepts a proton.
In the reaction: H2O + H2O <-> H3O+ + OH-, we can see that one of the water molecules (H2O) donates a proton, resulting in the formation of a hydronium ion (H3O+). Therefore, the water molecule acting as the proton donor (acid) is the one that loses a proton. On the other hand, the water molecule that accepts the proton (H+) becomes a hydroxide ion (OH-) and is, therefore, the proton acceptor (base).
So, in this reaction, we have the following conjugate acid-base pairs:
- Conjugate acid: H2O (donates a proton, becoming H3O+)
- Conjugate base: H2O (accepts a proton, becoming OH-)