# Can someone explain the Bronsted-Lowry definition of acids and bases?

## I get that the acid donates an H ion and the base excepts. However, it says, "To recognize a Bronsted-Lowry acid, look for the H at the beginning of the formula. Then it says, "To recognize a Bronsted-Lowry base, first make sure it is not an acid. Then, look for the N in the formula. Where does the N come in? I thought it was just the acid donating an H and a base accepting the H?

Feb 17, 2016

$\text{Acid + base " rightleftharpoons" Conjugate acid + conjugate base}$

#### Explanation:

Bronsted and Lowry characterized acid base behaviour on the basis of proton transfer, ${H}^{+}$. The acid was the proton donor, and the base the proton acceptor. I think your question refers to ammonia, i.e. $N {H}_{3}$, which is certainly a PROTON acceptor in aqueous solution, i.e.,

$N {H}_{3} \left(a q\right) + {H}^{+} \rightarrow N {H}_{4}^{+} + H {O}^{-}$

Capisce? So, in relation to your question, nitrogen containing species are often basic, and contain a neutral nitrogen ATOM, conceived to have a lone pair, available for donation to ${H}^{+}$.

Lewis proposed an alternative acid-base theory, in which the the base was the elecron pair donor, and the acid the electron pair acceptor, which allows a much broader church of acid-base behaviour. ${H}^{+}$, whatever this is, and IS very much a conceptual species, is a potent electron pair acceptor, i.e.

H^+ + ""^(-)OH rarr H_2O