Question

(a) What is the difference between the Arrhenius and the Bronsted-Lowry definitions of a base? (b) When ammonia is dissolved in water, it behaves both as an Arrhenius base and as a Bronsted-Lowry base. Explain.

Short answer

(a) The Arrhenius definition of a base is a substance that increases the concentration of hydroxide ions (\(OH^-\)) in water, while the Bronsted-Lowry definition of a base is a substance that can accept a proton (\(H^+\)) from another substance. The Arrhenius definition is limited to aqueous solutions whereas the Bronsted-Lowry definition is more widely accepted because it can be applied to a broader range of chemical reactions. (b) When ammonia (\(NH_3\)) is dissolved in water, it forms ammonium ions (\(NH_4^+\)) and hydroxide ions (\(OH^-\)) through the reaction: \(NH_3 (aq) + H_2O (l) \rightleftharpoons NH_4^+ (aq) + OH^- (aq)\). Ammonia behaves as an Arrhenius base by increasing the concentration of hydroxide ions (\(OH^-\)) and as a Bronsted-Lowry base by accepting protons (\(H^+\)) from water molecules.

Step-by-step solution

1. Understand the Arrhenius Definition of a Base

The Arrhenius definition of a base is a substance that increases the concentration of hydroxide ions (\(OH^-\)) when dissolved in water. In other words, an Arrhenius base is a substance that donates \(OH^-\) ions to the solution.

2. Understand the Bronsted-Lowry Definition of a Base

The Bronsted-Lowry definition of a base is a substance that can accept a proton (also known as a hydrogen ion, \(H^+\)) from another substance. In other words, a Bronsted-Lowry base is a proton acceptor.

3. Comparing the Two Definitions

The main difference between the Arrhenius and Bronsted-Lowry definitions is that the Arrhenius definition of a base is limited to aqueous solutions and focuses on the production of hydroxide ions, while the Bronsted-Lowry definition is more general and applies to any substance that can accept a proton. The Bronsted-Lowry definition is more widely accepted because it can be applied to a broader range of chemical reactions.

4. The Behavior of Ammonia in Water

When ammonia (\(NH_3\)) is dissolved in water, it reacts with water molecules to form ammonium ions (\(NH_4^+\)) and hydroxide ions (\(OH^-\)): \[ NH_3 (aq) + H_2O (l) \rightleftharpoons NH_4^+ (aq) + OH^- (aq) \]

5. Ammonia as an Arrhenius Base

As the equation in step 4 shows, when ammonia dissolves in water, it increases the concentration of hydroxide ions (\(OH^-\)). This behavior is consistent with the Arrhenius definition of a base, as it donates \(OH^-\) ions to the solution.

6. Ammonia as a Bronsted-Lowry Base

In the same equation from step 4, ammonia (\(NH_3\)) accepts a proton (\(H^+\)) from water to form ammonium ions (\(NH_4^+\)). This behavior is consistent with the Bronsted-Lowry definition of a base, as ammonia acts as a proton acceptor.

Conclusion

The difference between the Arrhenius and Bronsted-Lowry definitions of a base is that the Arrhenius definition is limited to substances that produce hydroxide ions in water, while the Bronsted-Lowry definition is more general and can apply to any substance that can accept a proton. When ammonia is dissolved in water, it behaves both as an Arrhenius base by increasing the concentration of hydroxide ions and as a Bronsted-Lowry base by accepting protons from water molecules.

Key concepts

Arrhenius Base

The concept of an Arrhenius base revolves simply around its ability to increase the number of hydroxide ions (\(OH^{-}\)) in an aqueous solution. According to Swedish scientist Svante Arrhenius, a base is a substance that when dissolved in water, dissociates to produce these negative ions. This is pivotal because hydroxide ions play a major role in creating the characteristic properties of a basic or alkaline solution, such as a bitter taste, slippery feel, and the ability to turn red litmus paper blue.

For instance, when sodium hydroxide (NaOH) is dissolved in water, it separates into sodium (\(Na^+\)) and hydroxide ions (\(OH^{-}\)), which causes the solution to become basic. It's essential to recognize that the Arrhenius definition strictly applies to reactions in aqueous solutions, which has limited its application in comparison to other definitions.

Bronsted-Lowry Base

Moving beyond the limitation of aqueous solutions comes the Bronsted-Lowry definition of a base. Here, a base is characterized as a proton acceptor. A proton (\(H^+\)), which is essentially a hydrogen atom stripped off its electron, can be donated by one substance and accepted by another. This broader concept introduced by chemists Johannes Nicolaus Brønsted and Thomas Martin Lowry, allows us to view not just hydroxide-ion-producing substances but any capable proton acceptor as a base, regardless of the solvent involved.

This definition greatly expands our perspective, enabling us to analyze a wider variety of chemical environments and reactions, even those taking place in a gaseous state or involving organic solvents. Additionally, the idea of proton exchange facilitates a deeper understanding of the dynamic, reversible nature of many acid-base interactions.

Chemical Reaction

Chemical reactions are processes in which one set of chemicals transforms into a new set of chemicals through the breaking and forming of bonds. Reactions often involve the rearrangement of atoms and changes in energy while obeying the law of conservation of mass. Within the context of acid-base reactions, like the reaction between an Arrhenius or Bronsted-Lowry base with water, we witness interactions that result in the generation of new ions or molecules.

Students can better grasp these interactions by observing that in acid-base reactions, there's a transfer of protons — acids donate protons while bases accept them. Hence, the reaction isn't merely about mixing substances; it's an orchestrated dance of particles that follow rules, such as electron configuration and energy minimization, to achieve a more stable arrangement.

Ammonia in Water

The dissolving of ammonia in water serves as an excellent example of these concepts at play. When ammonia (\(NH_3\)) combines with water (\(H_2O\)), it readily accepts a proton from water to form ammonium (\(NH_4^+\)) and hydroxide ions (\(OH^{-}\)), demonstrating ammonia's versatile behavior as both an Arrhenius and a Bronsted-Lowry base. In this chemical dance, ammonia's need for an extra proton (as it has an electron pair ready to bond) and water's ability to donate a proton (being protic and polar) meet harmoniously.

This reaction also illustrates the essential principle of acid-base chemistry: the interplay and balance of proton donators and acceptors. By understanding that bases can increase hydroxide ion concentration and also accept protons, students can better navigate through the complexities of various chemical systems and their behaviors.