Question

(a) What is the difference between the Arthenius 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

The main difference between the Arrhenius and Brønsted-Lowry definitions of a base is that the Arrhenius definition focuses on the donation of hydroxide ions (\(OH^-\)) in aqueous solutions, while the Brønsted-Lowry definition focuses on proton acceptance, regardless of the medium. When ammonia (NH\(_3\)) is dissolved in water, it increases the hydroxide ion concentration and accepts protons from water molecules, making it both an Arrhenius and a Brønsted-Lowry base.

Step-by-step solution

Define Arrhenius base

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

Define Brønsted-Lowry base

The Brønsted-Lowry definition states that a base is a substance that can accept a proton (H\(^+\)) or, in other words, a substance that acts as a proton acceptor. In this case, the concept of a base is not limited to aqueous solutions, as with the Arrhenius definition.

Compare both definitions

The main difference between the Arrhenius and Brønsted-Lowry definitions of a base is that the Arrhenius definition only applies to aqueous solutions and focuses on the donation of \(OH^-\) ions, while the Brønsted-Lowry definition applies to bases that can accept protons, regardless of the medium.

Explain how ammonia is an Arrhenius base

When ammonia (NH\(_3\)) is dissolved in water, it reacts with water molecules to form ammonium ions (NH\(_4^+\)) and hydroxide ions (\(OH^-\)). The chemical equation for this reaction is: \[NH_3 (aq) + H_2O (l) \rightleftharpoons NH_4^+ (aq) + OH^- (aq)\] By increasing the concentration of hydroxide ions in the aqueous solution, ammonia acts as an Arrhenius base.

Explain how ammonia is a Brønsted-Lowry base

In the same reaction between ammonia and water, we can observe that ammonia accepts a proton from the water molecule, forming an ammonium ion (NH\(_4^+\)). This behavior of accepting a proton from a donor substance classifies ammonia as a Brønsted-Lowry base. In summary, when ammonia is dissolved in water, it increases the hydroxide ion concentration, making it an Arrhenius base, and it also accepts protons from water molecules, making it a Brønsted-Lowry base.

Key concepts

Arrhenius Base

In the world of chemistry, one of the first theories students encounter is the Arrhenius theory, which provides a fundamental understanding of acid-base chemistry. An Arrhenius base is characterized by its ability to increase the concentration of hydroxide ions (\(OH^-\)) in a solution. This increase occurs when the base disassociates in water, releasing hydroxide ions as part of the process. When these ions are released, they can react with hydrogen ions (\(H^+\)) to form water, thus reducing the overall acidity of the solution.

Consider sodium hydroxide (\(NaOH\)), a classic example. When dissolved in water, it dissociates completely to give sodium (\(Na^+\)) and hydroxide (\(OH^-\)) ions:
\begin{align*}NaOH (s) &\rightarrow Na^+ (aq) + OH^- (aq)\because{align*}This behavior typifies the Arrhenius concept of a base and has great practical implications in understanding how substances neutralize acids and affect pH levels in various chemical processes.

Bronsted-Lowry Base

Moving beyond the Arrhenius theory, the Bronsted-Lowry theory broadens the horizon to include more diverse chemical environments beyond aqueous solutions. A Bronsted-Lowry base is defined not just by its ability to produce hydroxide ions but by its capability to accept protons (\(H^+\)). This more inclusive definition allows us to identify bases in a variety of reactions, including those with no hydroxide ions present.

For example, ammonia (\(NH_3\) acts as a Bronsted-Lowry base when it accepts a proton from a water molecule, producing the ammonium ion (\(NH_4^+\)):\begin{align*}NH_3 (aq) + H_2O (l) &\rightleftharpoons NH_4^+ (aq) + OH^- (aq)\because{align*}This versatility provided by the Bronsted-Lowry theory allows chemists to identify a broader range of substances as bases, facilitating a deeper understanding of complex chemical reactions.

Ammonia as a Base

Ammonia (\(NH_3\)), a colorless gas with a pungent smell, is a fascinating substance that acts as a base in both Arrhenius and Bronsted-Lowry frameworks. When ammonia is dissolved in water, it undergoes a reversible reaction, where it accepts a proton from water (\(H_2O\)) and forms ammonium (\(NH_4^+\)) and hydroxide (\(OH^-\)) ions:
\begin{align*}NH_3 (aq) + H_2O (l) &\rightleftharpoons NH_4^+ (aq) + OH^- (aq)\because{align*}In this reaction, the creation of hydroxide ions signifies its role as an Arrhenius base. Simultaneously, the acceptance of a proton classifies ammonia as a Bronsted-Lowry base. This dual nature showcases the versatility of ammonia and its significant role in acid-base chemistry, particularly in the formation of many ammonium compounds.

Proton Acceptors

The concept of proton acceptors is integral to the Bronsted-Lowry definition of a base. Substances that are proton acceptors are able to accept a hydrogen ion (\(H^+\)), which is simply a proton, from another substance. The substance donating the proton is typically referred to as the acid. This interaction between proton donors and acceptors is what constitutes an acid-base reaction in the Bronsted-Lowry sense.

It's interesting to note that water (\(H_2O\)) can act as both a proton donor and a proton acceptor, making it amphiprotic. This characteristic of water plays a crucial role in many chemical reactions, including neutralization and hydrolysis. Understanding the role of proton acceptors helps students appreciate the dynamic nature of chemical reactions and the subtleties of acid-base interactions.