In the following acid-base reaction, how would you identify the acid, base, and their conjugate acids and bases: NH4+ + HCO3- --> NH3 + H2CO3?

A useful way of thinking about this is:

• Seeing which ion donates a proton (as a Brønsted-Lowry acid) by accepting a pair of electrons (as a Lewis acid)
• Seeing which ion accepts the proton (as a Brønsted-Lowry base) by donating a pair of electrons (as a Lewis base).

Another way to say it is in general, an acid has a proton it can donate, and a base wants a proton.

The change from #"NH"_4^(+)# to #"NH"_3# implies a proton was donated from #"NH"_4^(+)#, so simply put, #"NH"_4^(+)# is the acid, and #"NH"_3# is the base. Since #"NH"_3# is the neutral state, #"NH"_4^(+)# is the conjugate acid.

The change from #"HCO"_3^(-)# to #"H"_2"CO"_3# implies a proton was accepted by #"HCO"_3^(-)#, so simply put, #"HCO"_3^(-)# is the base, and #"H"_2"CO"_3# is the acid. Since #"H"_2"CO"_3# is the neutral state, #"HCO"_3^(-)# is the conjugate base.

Overall, we have:

#"NH"_4^(+)#

• Proton donor (Brønsted-Lowry acid)
• Electron acceptor (Lewis acid)

#"NH"_3#

• Proton acceptor (Brønsted-Lowry base)
• Electron donor (Lewis base)

#"HCO"_3^(-)#

• Proton acceptor (Brønsted-Lowry base)
• Electron donor (Lewis base)

#"H"_2"CO"_3#

• Proton donor (Brønsted-Lowry acid)
• Electron acceptor (Lewis acid)

Visually, this is how it is represented: