Question

(a) What is the difference between the Arrhenius and the Brønsted-Lowry definitions of an acid? (b) \(\mathrm{NH}_{3}(g)\) and \(\mathrm{HCl}(\mathrm{g})\) react to form the ionic solid \(\mathrm{NH}_{4} \mathrm{Cl}(s)\). Which substance is the Bronsted-Lowry acid in this reaction? Which is the Bronsted-Lowry base?

Short answer

(a) The main difference between the Arrhenius and Brønsted-Lowry definitions of an acid is that the Arrhenius definition focuses on the increase of hydrogen ions (\(\mathrm{H}^{+}\)) in water, while the Brønsted-Lowry definition focuses on the ability to donate a proton (\(\mathrm{H}^{+}\)) to another substance, making the latter more general. (b) In the reaction \(\mathrm{NH}_{3}(g) + \mathrm{HCl}(g) \rightarrow \mathrm{NH}_{4} \mathrm{Cl}(s)\), \(\mathrm{HCl}(\mathrm{g})\) is the Brønsted-Lowry acid, as it donates a proton, and \(\mathrm{NH}_{3}(g)\) is the Brønsted-Lowry base, as it accepts a proton.

Step-by-step solution

(a) Comparing Acid Definitions

In order to differentiate between the Arrhenius and the Brønsted-Lowry definitions of an acid, we need to understand both definitions: 1. Arrhenius Definition: According to the Arrhenius definition, an acid is a substance that increases the concentration of hydrogen ions (\(\mathrm{H}^{+}\)) when dissolved in water, while a base is a substance that increases the concentration of hydroxide ions (\(\mathrm{OH}^{-}\)) when dissolved in water. 2. Brønsted-Lowry Definition: According to the Brønsted-Lowry definition, an acid is a substance that donates a proton (\(\mathrm{H}^{+}\)) to another substance, while a base is a substance that accepts a proton from another substance. The main difference between these two definitions is that the Arrhenius definition is limited to acids and bases in aqueous solutions, while the Brønsted-Lowry definition is more general and can be applied to various substances and reactions, including those that do not involve water.

(b) Identifying the Brønsted-Lowry Acid and Base

First, let's write down the given reaction: \(\mathrm{NH}_{3}(g) + \mathrm{HCl}(g) \rightarrow \mathrm{NH}_{4} \mathrm{Cl}(s) \) Now, we need to identify which substance is donating a proton and which one is accepting it. To do this, we can rewrite the formation of \(\mathrm{NH}_{4} \mathrm{Cl}(s)\) as the following proton transfer reaction: \(\mathrm{NH}_{3}(g) + \mathrm{H}^{+} \rightarrow \mathrm{NH}_{4}^{+} \) \(\mathrm{Cl}^{-} \rightarrow \mathrm{HCl}(g) \) From these reactions, we can see that \(\mathrm{NH}_{3}(g)\) is accepting a proton to become \(\mathrm{NH}_{4}^{+}\), while \(\mathrm{HCl}(\mathrm{g})\) is donating a proton to form \(\mathrm{Cl}^{-}\). Therefore, according to the Brønsted-Lowry definition: - The acid is \(\mathrm{HCl}(\mathrm{g})\), as it donates a proton. - The base is \(\mathrm{NH}_{3}(g)\), as it accepts a proton.

Key concepts

Arrhenius definition

The Arrhenius definition of acids and bases focuses on their behavior in water. According to this definition, an acid is a substance that increases the concentration of hydrogen ions \( (\mathrm{H}^+) \) when it is dissolved in water. In simpler terms, an Arrhenius acid releases hydrogen ions into the water, making the solution more acidic. On the flip side, an Arrhenius base increases the concentration of hydroxide ions \( (\mathrm{OH}^-) \) in water, making the solution more basic.

While useful, this definition is limited because it only considers reactions in aqueous solutions. The Arrhenius definition does not apply to chemical processes where water is not present, nor does it account for the role of the solvent in reactions. Despite these limitations, this definition laid the groundwork for our understanding of acids and bases in chemistry.

Brønsted-Lowry definition

The Brønsted-Lowry definition expands upon the limitations of the Arrhenius model by considering acids and bases in broader contexts. According to Brønsted-Lowry, an acid is a substance that donates a proton \( (\mathrm{H}^+) \) to another substance, while a base is one that accepts a proton.

This definition liberates acids and bases from the constraint of aqueous solutions, allowing for reactions in gas, non-aqueous solvents, or even solid phases. The `proton transfer` concept is central here, where the acid donates a hydrogen ion to a base.
An example of this reaction can be seen in the formation of ammonium chloride \( (\mathrm{NH}_4\mathrm{Cl}) \). Here, hydrogen chloride \( (\mathrm{HCl}) \) acts as a Brønsted-Lowry acid, donating a proton to ammonia \( (\mathrm{NH}_3) \), which acts as the base by accepting this proton. This versatility makes the Brønsted-Lowry definition highly applicable in a wide array of chemical scenarios.

Proton transfer reaction

Proton transfer reactions are a cornerstone of the Brønsted-Lowry acid-base theory and play a critical role in many chemical processes. In these reactions, a proton \( (\mathrm{H}^+) \) is transferred from an acid to a base. This transfer leads to the formation of a conjugate base (formerly the acid) and a conjugate acid (formerly the base).

In the reaction between ammonia \( (\mathrm{NH}_3) \) and hydrogen chloride \( (\mathrm{HCl}) \), \( \mathrm{HCl} \) donates a proton to \( \mathrm{NH}_3 \), forming ammonium \( (\mathrm{NH}_4^+) \) and chloride ions \( (\mathrm{Cl}^-) \).
These ions then combine to form ammonium chloride \( (\mathrm{NH}_4\mathrm{Cl}) \), a solid ionic compound. Such reactions are not only essential in academic chemistry but are also foundational to many industrial and biological processes. Understanding proton transfer helps students grasp a fundamental mechanism that facilitates countless reactions across different environments, whether in labs, nature, or daily life.