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: