How do you identify Bronsted acids and bases?

I once read a comment that stated - trying to classify a substance as a Bronsted acid or base is a little like trying to judge the character of a person. Is long as they are alone and isolated, you cannot tell. It is only through their interactions with other people that you see characteristics like generosity, etc.

Bronsted acids and bases are like that. If a molecule or ion is isolated from other compounds, you cannot tell whether it will be a proton donor or acceptor. It is only when it interacts with other molecules that you can identify its character.

Here are two examples:

#HSO_4^-# #+ NH_3 rarr NH_4^+ + SO_4^(2-)#

Here, the #HSO_4^-# ion donates a proton to #NH_3# and is a Bronsted acid.

#HCl + HSO_4^-# #rarr H_2SO_4 + Cl^-#

This time, #HSO_4^-# accepts a proton from #HCl# and is a Bronsted base.

So, which is it, acid or base? It is both, or better - it can be either, depending on what you mix it with.

A #"Bronsted acid"# is a species that will donate a proton to the solvent (which is typically water). And thus, we can write the generalized acid-base reaction:

#HA+H_2O(l) rightleftharpoonsH_3O^(+) + A^(-)#

As in all chemical reactions, both charge and mass must be conserved. What does this mean? Is it conserved here? The macroscopic observable for this reaction is typically a marked change in the #pH# of the solution (in which direction?), or more rarely, a marked increase in electrical conductivity.

And thus Bronsted acids are conceived to donate a proton to the water solvent, to form the conjugate acid, here hydronium ion, #H_3O^+#, and some conjugate base, #A^-#, or what have you.

What are the conjugate bases of common acids such as #HNO_3#, #H_2SO_4#, and #H_3PO_4#? What is the conjugate base of #HSO_4^-#?