Identify the acid and the base in the following reactions: (a) \(\mathrm{CH}_{3} \mathrm{NH}_{2}(\mathrm{aq})+\mathrm{H}_{3} \mathrm{O}^{+}(\mathrm{aq}) \rightarrow\) \(\mathrm{CH}_{3} \mathrm{NH}_{3}{ }^{*}(\mathrm{aq})+\mathrm{H}_{2} \mathrm{O}(\mathrm{l})\) (b) \(\mathrm{C}_{2} \mathrm{H}_{3} \mathrm{NH}_{2}\) (aq) \(+\mathrm{HCl}\) (aq) \(\rightarrow\) \(\mathrm{C}_{2} \mathrm{H}_{5} \mathrm{NH}_{3}{ }^{+}(\mathrm{aq})+\mathrm{Cl}^{-}(\mathrm{aq})\) (c) \(\mathrm{CaO}\) (s) \(+2 \mathrm{HI}(\mathrm{aq}) \rightarrow \mathrm{CaI}_{2}\) (aq) \(+\mathrm{H}_{2} \mathrm{O}\) (l)

Understanding the Bronsted-Lowry theory is fundamental in identifying acids and bases in a chemical reaction. This theory extends the definition of acids and bases beyond the classical idea which limited acids and bases to substances that produce hydrogen or hydroxide ions in water.

The theory states that an acid is a substance capable of donating a proton (a hydrogen ion, H+), while a base is a substance capable of accepting a proton. This proton-exchange perspective allows us to classify substances as acids or bases based on their behavior in a reaction, even if they are not in an aqueous solution.

One crucial aspect of the Bronsted-Lowry theory is that it introduced the concept of conjugate acid-base pairs. In a chemical reaction, when an acid donates a proton, it becomes a base, known as the conjugate base. Similarly, when a base accepts a proton, it becomes an acid, referred to as the conjugate acid. This duality is significant because it shows the reversibility of acid-base reactions and explains how substances can act as either acids or bases depending on the context of the reaction.

The Bronsted-Lowry theory relies on the roles of proton donors and acceptors to categorize substances. A proton donor, which we call an acid, has a detachable hydrogen ion within its molecular structure that it can release. On the other hand, a proton acceptor, called a base, has the capacity to bind with these free hydrogen ions.

Key Example

In the reaction (a), \(\mathrm{H}_3\mathrm{O}^+\) donates a proton to \(\mathrm{CH}_3\mathrm{NH}_2\), making \(\mathrm{H}_3\mathrm{O}^+\) the proton donor (acid), and \(\mathrm{CH}_3\mathrm{NH}_2\) the proton acceptor (base). The ability to donate or accept protons is not a fixed property but context-dependent, as substances can act differently in the presence of other chemical species.

When analyzing chemical equations, spotting the proton transactions is critical: a decrease in hydrogen atoms in one species indicates it's acting as an acid, whereas an increase suggests base behavior. This transfer is evident in acid-base reactions and is a reliable indicator to identify the role of each reactant.

Acid-base reactions, also known as neutralization reactions, are processes in which an acid and a base react to form water and a salt. However, from the Bronsted-Lowry perspective, the essential feature of these reactions is the transfer of protons from the acid to the base.

The equations provided in the exercise are all examples of acid-base reactions. For instance, in reaction (b), \(\mathrm{HCl}\) donates a proton to \(\mathrm{C}_2\mathrm{H}_3\mathrm{NH}_2\), resulting in the formation of \(\mathrm{C}_2\mathrm{H}_5\mathrm{NH}_3^+\) and \(\mathrm{Cl}^-\), illustrating how the acid (hydrochloric acid) reacts with the base (amine) to produce a salt and modifies the base into its conjugate acid.

Moreover, these reactions are equilibrium processes. This means they can occur in both directions, a point often overlooked. In a given reaction, the acid and base on one side of the equation will form conjugates on the opposite side. If conditions change or if the conjugates are strong enough, the reaction can reverse, demonstrating the dynamic nature of acid-base chemistry and emphasizing the importance of understanding how substances can interchange their roles as acids or bases.