What are the common definitions of an acid? How can #"H"_2"CO"_3# and #"HCO"_3^-# both be acids if #"HCO"_3^-# is a conjugate base, and what does it mean for the #"pH"# of a #"HCO"_3^-# solution??
1 Answer
Bronsted-Lowry definition:
It is an acid if it donates a proton to another compound.
Lewis definition:
It is an acid if it accepts electrons from another compound.
Arrhenius definition:
It is an acid if it donates a proton to the solution upon dissociation.
None of these acid definitions has to do with the resultant pH or pOH. We calculate them after-the-fact.
These acid definitions pertain only to the compound's function in relationship to another compound, or the compound's function after dissociation.
So,
It does all three things. It donates a proton to another compound, it accepts electrons from another compound (to donate a proton), and it dissociates to give a proton to solution.
Also, and it might seem strange to you, but
It can donate a proton to another compound, accept electrons from another compound (to donate its proton), donate a proton to solution upon dissociation, donate electrons (any of the oxygens), AND accept a proton (by donating electrons). These, respectively, correspond to the acid and base definitions used in the previous paragraph.
In terms of pKa, we can alternatively define acidity as an extent of dissociation, based on the Bronsted-Lowry definition, relative to the solvent, or another compound.
Technically, when an acid is put into water, it will fit both the Bronsted-Lowry and Arrhenius definition because water is both a solvent AND the compound that gets a proton. When this happens, we can look at the pKa.
The pKa
