# What is the conjugate acid and base of HSO_4^-?

Jun 27, 2016

Here's what I got.

#### Explanation:

The thing to remember about conjugate acids is that they are the chemical species that is formed when a Bronsted-Lowry base accepts one proton, ${\text{H}}^{+}$.

This means that in order to find the conjugate acid of a substance that can act as a Bronsted-Lowry base, all you have to do is add a proton to it.

But keep in mind that a proton carries a $1 +$ charge, so make sure that you take this into account.

In your case, the hydrogen sulfate anion, ${\text{HSO}}_{4}^{-}$, can act as a Bronsted-Lowry base and accept a proton to form sulfuric acid, ${\text{H"_2"SO}}_{4}$

overbrace("HSO"_ (4(aq))^(-))^(color(blue)("base")) + "H"_ ((aq))^(+) -> overbrace("H"_ 2"SO" _(4(aq)))^(color(darkgreen)("conjugate acid"))

Notice that because you're adding a $1 +$ charge to a compound that has a $1 -$ charge, the resulting compound will have a zero net charge.

Similarly, conjugate bases are chemical species that are formed when a Bronsted-Lowry acid donates one proton.

This means that you can find the conjugate base of a Bronsted-Lowry acid by removing a proton from it.

In your case, the hydrogen sulfate anion can act as a Bronsted-Lowry acid and donate a proton to form the sulfate anion, ${\text{SO}}_{4}^{2 -}$.

overbrace("HSO"_ (4(aq))^(-))^(color(darkgreen)("acid")) -> "H"_ ((aq))^(+) + overbrace("SO"_ (4(aq))^(2-))^(color(blue)("conjugate base"))

Once again, notice that the charge is balanced because removing a $1 +$ charge from a compound that has a $1 -$ charge will give you a compound that has a $2 -$ net charge.