Question

The conjugate base of $\text{H}_{2}\text{PO}_{4}^{-}$ is

• A. $\text{HPO}_{4}^{2-}$
• B. $\text{PO}_{4}^{3-}$
• C. $\text{H}_{3}\text{PO}_{4}$
• D. $\text{H}_{2}\text{PO}_{4}^{2-}$

Hint:

Acid minus $\text{H}^{+}$ $\rightarrow$ Conjugate Base. Conversely, adding an $\text{H}^{+}$ proton to a base gives you its conjugate acid.

Answer 1

The Correct Option is A

Step 1: Recall the meaning of conjugate base.
In the Bronsted-Lowry idea, an acid gives away one proton, which is one $H^+$. What is left behind is called its conjugate base.

Step 2: Write the simple rule.
To get the conjugate base, remove one $H$ from the formula and make the charge more negative by one.

Step 3: Start with the given species.
The given acid is $H_2PO_4^-$. It has two hydrogens and a charge of minus one.

Step 4: Remove one hydrogen.
Taking one $H$ away from $H_2PO_4^-$ leaves the $HPO_4$ part.

Step 5: Adjust the charge.
Removing $H^+$ makes the charge go from minus one to minus two. So we get $HPO_4^{2-}$.

Step 6: State the answer.
So the conjugate base of $H_2PO_4^-$ is the hydrogen phosphate ion. \[ \boxed{HPO_4^{2-}} \]