Question

Which of the following is both a Bronsted acid and a Bronsted base in water? A. \(\mathrm{H}_{2} \mathrm{PO}_{4}^{-}\) B. \(\mathrm{H}_{2} \mathrm{CO}_{3}\) C. \(\mathrm{NH}_{3}\) D. \(\mathrm{NH}_{4}^{+}\) E. \(\mathrm{Cl}^{-}\)

Short answer

Answer: Option A: \(\mathrm{H}_{2} \mathrm{PO}_{4}^{-}\) acts as both a Bronsted acid and a Bronsted base in water.

Step-by-step solution

Option A: \(\mathrm{H}_{2} \mathrm{PO}_{4}^{-}\)

In water, \(\mathrm{H}_{2} \mathrm{PO}_{4}^{-}\) can donate a proton to become \(\mathrm{HPO}_{4}^{2-}\). It can also accept a proton to become \(\mathrm{H}_{3} \mathrm{PO}_{4}\). Since it can both donate and accept a proton, it can act as both a Bronsted acid and a Bronsted base.

Option B: \(\mathrm{H}_{2} \mathrm{CO}_{3}\)

In water, \(\mathrm{H}_{2} \mathrm{CO}_{3}\) can donate a proton to become \(\mathrm{HCO}_{3}^{-}\), but it cannot accept a proton to become \(\mathrm{H}_{3}\mathrm{CO}_{3}^{+}\), which does not exist in the aqueous solution. Hence, it can only act as a Bronsted acid.

Option C: \(\mathrm{NH}_{3}\)

In water, \(\mathrm{NH}_{3}\) can accept a proton to become \(\mathrm{NH}_{4}^{+}\), but it cannot donate a proton as it has no hydrogen ions to donate. Hence, it can only act as a Bronsted base.

Option D: \(\mathrm{NH}_{4}^{+}\)

In water, \(\mathrm{NH}_{4}^{+}\) can donate a proton to become \(\mathrm{NH}_{3}\), but it cannot accept a proton as it has no lone pairs to accept a proton. Hence, it can only act as a Bronsted acid.

Option E: \(\mathrm{Cl}^{-}\)

In water, \(\mathrm{Cl}^{-}\) can accept a proton to become \(\mathrm{HCl}\), but it cannot donate a proton as it has no hydrogen ions to donate. Hence, it can only act as a Bronsted base. So, the correct answer is:

Answer:

Option A: \(\mathrm{H}_{2} \mathrm{PO}_{4}^{-}\) is the substance that acts as both a Bronsted acid and Bronsted base in water.