Question

Which of the following cannot behave like both bronsted acid as well as bronsted base? (A) \(\mathrm{H}_{3} \mathrm{O}^{\oplus}\) (B) \(\mathrm{H}_{2} \mathrm{O}\) (C) \(\mathrm{NH}_{3}\) (D) \(\mathrm{HSO}_{4}^{-}\)

Short answer

The answer is (A) \(\mathrm{H}_{3} \mathrm{O}^{\oplus}\), which cannot behave like both a Bronsted acid and a Bronsted base because it does not have a lone pair of electrons to accept a proton, and thus can only act as a Bronsted-Lowry acid.

Step-by-step solution

Analysis of Species A: \(\mathrm{H}_{3} \mathrm{O}^{\oplus}\)

The \(\mathrm{H}_{3} \mathrm{O}^{\oplus}\) ion, also known as hydronium ion, has three hydrogens that can be donated, but it does not have a lone pair of electrons to accept any additional proton. Therefore, it can act as a Bronsted-Lowry acid but not as a Bronsted-Lowry base.

Analysis of Species B: \(\mathrm{H}_{2} \mathrm{O}\)

Water (\(\mathrm{H}_{2} \mathrm{O}\)) has two hydrogens that it can donate and also two lone pairs of electrons to accept a proton. Hence, it can act as both a Bronsted-Lowry acid and a Bronsted-Lowry base.

Analysis of Species C: \(\mathrm{NH}_{3}\)

Ammonia (\(\mathrm{NH}_{3}\)) does not have any hydrogen that it can donate but it has a lone pair of electrons, which allows it to accept a proton. Hence, it can act as a Bronsted-Lowry base but not as a Bronsted-Lowry acid.

Analysis of Species D: \(\mathrm{HSO}_{4}^{-}\)

The \(\mathrm{HSO}_{4}^{-}\) ion (the hydrogen sulfate ion) does have a hydrogen that it can donate, and the sulfur and oxygen atoms both have lone pairs of electrons to accept a proton. Therefore, it can act as both a Bronsted-Lowry acid and a Bronsted-Lowry base.

Summary and answer to the exercise

Therefore, in light of the analysis, the answer is (A) \(\mathrm{H}_{3} \mathrm{O}^{\oplus}\), which cannot behave like both a Bronsted acid as well as a Bronsted base because it does not have a lone pair of electrons to accept a proton, and thus can only act as a Bronsted-Lowry acid.

Key concepts

Hydronium Ion

The hydronium ion, represented as \text{(A)} \(\mathrm{H}_{3} \mathrm{O}^{\oplus}\), plays a critical role in acid-base chemistry. As a positive ion (\text{cation}), it is formed when a water molecule gains a proton (\text{H}^+). Hydronium is considered an acid in the Bronsted-Lowry definition, which defines an acid as a proton donor.

However, when it comes to being a base, the hydronium ion falls short because it lacks lone pairs of electrons on oxygen to accept another proton. This is why in our exercise \text{(A)} \(\mathrm{H}_{3} \mathrm{O}^{\oplus}\) could not act as a Bronsted base. Understanding the hydronium ion is key to comprehending pH levels, as it directly affects the acidity of solutions. Substances with a higher concentration of hydronium ions are more acidic.

Acid-Base Reactions

Acid-base reactions are fundamental to chemistry and involve the exchange of protons between molecules. According to the Bronsted-Lowry theory, an acid is a substance that can donate a proton (\text{H}^+), while a base is a substance that can accept a proton.

In our exercise, analyzing the reaction potential of different molecules, we considered their ability to both donate and accept protons. Water (\(\mathrm{H}_{2} \mathrm{O}\)) and the hydrogen sulfate ion (\(\mathrm{HSO}_{4}^{-}\)) can act as both acids and bases because they have the necessary components: hydrogen atoms to donate and lone pairs to accept protons. Ammonia (\(\mathrm{NH}_{3}\)), with its lone pair, acts only as a base. The crux of these reactions is the transfer of protons, influencing the reactivity and properties of the substances involved.

Lone Pair of Electrons

A lone pair of electrons refers to a pair of electrons associated with an atom that is not shared or bonded with another atom. They are important for understanding the reactivity of molecules, especially in the case of bases in Bronsted-Lowry acid-base theory.

For an entity to act as a base, as seen in molecule \(\mathrm{NH}_{3}\), these non-bonding electrons are capable of binding to a proton from an acid. In contrast, the hydronium ion lacks this electron pair, rendering it incapable of such action. Molecules or ions with lone pairs can participate in a variety of chemical reactions and their presence can significantly alter the physical properties, such as the boiling point or solubility, of a substance.