# For HSO_4^(-), NH_3(aq), and NH_4^(+) which species are Bronsted acids, and which are Bronsted bases?

May 11, 2017

As far as I can see (and you made me pull out my spex), I would amend $c$, $e$, and $f$.....Why?

#### Explanation:

$c .$, you got $\text{bisulfate ion}$; in aqueous solution this is a moderately strong acid. Certainly we treat ${H}_{2} S {O}_{4}$ as a diacid in aqueous solution.

$e .$, you got $\text{ammonium ion}$; in aqueous solution this is a weak Bronsted acid:

$N {H}_{4}^{+} + {H}_{2} O \left(l\right) r i g h t \le f t h a r p \infty n s {H}_{3} {O}^{+} + N {H}_{3} \left(a q\right)$

$f .$, you got $\text{ammonia}$; in aqueous solution this is a weak Bronsted base, that gives equilibrium quantities of hydroxide at equilibrium, i.e.:

$N {H}_{3} \left(a q\right) + {H}_{2} O r i g h t \le f t h a r p \infty n s N {H}_{4}^{+} + H {O}^{-}$

Note that I have of course assumed an aqueous solution for each of the acids and bases.

May 11, 2017

Everything looks perfect, except for (e) and (f)

#### Explanation:

You've pretty much nailed everything, but $N {H}_{3}$ and $N {H}_{4}^{+}$ are not good at being an acid and base respectively.

In theory, they could probably donate/receive that extra proton, but the compounds they produce ($N {H}_{2}^{-}$ and $N {H}_{5}^{2 +}$) are not very stable. Also, as you do more labs and other work on the topic of acids & bases, you'll see that you never really consider either of these compounds as being both an acid and base.

Hence, I'd say that $N {H}_{3}$ is just a Brønsted base, and $N {H}_{4}$ is just a Brønsted acid.

Hope that helps :)