Question

How are the Arrhenius and Bronsted-Lowry acid-base definitions different? How are they similar? Name two BronstedLowry bases that are not Arrhenius bases. Can you do the same for acids? Explain.

Short answer

Arrhenius acids/bases involve H⁺/OH⁻ ions in water; Bronsted-Lowry involves proton transfers in any environment. Examples of Bronsted-Lowry bases not Arrhenius: NH₃, CH₃COO⁻.

Step-by-step solution

Understand the Arrhenius Definition

The Arrhenius definition states that an acid is a substance that increases the concentration of hydrogen ions (H⁺) in aqueous solution, while a base is a substance that increases the concentration of hydroxide ions (OH⁻) in aqueous solution.

Understand the Bronsted-Lowry Definition

The Bronsted-Lowry definition describes acids as proton (H⁺) donors and bases as proton (H⁺) acceptors. This definition is not limited to aqueous solutions and can be applied to reactions in any solvent or even in the gas phase.

Identify the Differences

The key difference between the two definitions is their scope. The Arrhenius definition is limited to aqueous solutions and involves the production of H⁺ and OH⁻ ions, whereas the Bronsted-Lowry definition is broader and is based on the transfer of protons, applicable in various environments.

Identify the Similarities

Both definitions describe acids in terms of their interactions with hydrogen ions (protons). In the Arrhenius framework, acids increase H⁺ concentration in water, while in the Bronsted-Lowry framework, acids donate H⁺ ions.

Identify Bronsted-Lowry Bases That Are Not Arrhenius Bases

Examples of Bronsted-Lowry bases that are not Arrhenius bases include ammonia (NH₃) and the acetate ion (CH₃COO⁻). These substances can accept protons without producing OH⁻ ions in solution.

Identify If There Are Bronsted-Lowry Acids That Are Not Arrhenius Acids

All Bronsted-Lowry acids (proton donors) should fit within the Arrhenius definition since their donation of protons invariably increases H⁺ concentration in aqueous solution. Therefore, there might be no Bronsted-Lowry acids that are not Arrhenius acids.

Key concepts

Arrhenius acid-base definition

The Arrhenius acid-base definition is one of the earliest concepts in chemistry. It focuses on the behavior of substances in aqueous solutions.

Arrhenius is specific to aqueous solutions

According to Arrhenius, an acid is any substance that increases the concentration of hydrogen ions (H⁺) when dissolved in water. Conversely, a base increases the concentration of hydroxide ions (OH⁻) in water.
So, when an Arrhenius acid dissolves in water, it releases H⁺ ions. And when an Arrhenius base dissolves, it releases OH⁻ ions.

Bronsted-Lowry acid-base theory

The Bronsted-Lowry acid-base theory expands the concept of acids and bases beyond just those in aqueous solutions.
In this theory, an acid is defined as a proton (H⁺) donor, and a base is defined as a proton acceptor.

Proton donors and acceptors

The role of proton donors and acceptors is central to the Bronsted-Lowry theory.

• Acids are substances that donate H⁺ ions.
• Bases are substances that accept H⁺ ions.

This theory doesn't rely solely on the presence of water, so it can be applied to acid-base reactions in gas phases and non-aqueous solutions.

Acid-base reactions

Acid-base reactions can be summarized by the transfer of a proton from an acid to a base.
For example, in the reaction between ammonia (NH₃) and hydrogen chloride (HCl), the HCl donates a proton to the NH₃, forming ammonium (NH₄⁺) and chloride (Cl⁻) ions.
This highlights the flexibility of the Bronsted-Lowry definition.

Aqueous solutions versus non-aqueous environments

The Arrhenius definition is limited to environments where water is present. However, the Bronsted-Lowry theory is applicable in both aqueous and non-aqueous environments.
For instance, in liquid ammonia (a non-aqueous solvent), ammonia can act as a base by accepting a proton from an acid-like ammonium ion. This demonstrates the broader applicability of the Bronsted-Lowry theory.