Question

Why is the Arrhenius acid-base definition too limited? Give an example for which the Arrhenius definition does not apply.

Short answer

Arrhenius' definition is limited to aqueous solutions and does not apply to acid-base reactions in non-aqueous solutions or gases, such as NH₃ and HCl forming NH₄Cl.

Step-by-step solution

- Understanding Arrhenius Definition

Under Arrhenius' definition, 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.

- Identify the Limitation

The main limitation of Arrhenius' definition is that it only applies to substances in aqueous solutions and does not account for acid-base reactions that occur in non-aqueous solutions or gases.

- Provide an Example

Consider the reaction between ammonia (NH₃) and hydrogen chloride gas (HCl). This reaction occurs in the absence of water and forms ammonium chloride (NH₄Cl), where NH₃ acts as a base, accepting a proton from HCl, which acts as an acid. According to Arrhenius' definition, both NH₃ and HCl do not fit the definition as they do not produce OH⁻ or H⁺ in aqueous solution.

Key concepts

Understanding Acid-Base Reactions

Acid-base reactions are fundamental chemical processes that occur when acids and bases interact. According to the Arrhenius definition, an acid is a substance that increases the concentration of hydrogen ions (H⁺) in an aqueous solution. A base, on the other hand, increases the concentration of hydroxide ions (OH⁻) in an aqueous solution.
However, these reactions can occur in different environments beyond aqueous solutions. In essence, an acid-base reaction involves the transfer of protons (H⁺) from one substance to another. In many acid-base reactions, an acid will donate a proton, while a base will accept that proton.
Understanding the basic mechanism of acid-base reactions is critical in appreciating the limitations of the Arrhenius definition. By exploring other definitions, such as those by Brønsted-Lowry and Lewis, you can gain a comprehensive understanding of how these reactions function in various contexts.

Exploring Aqueous Solutions

Aqueous solutions are simply solutions where water acts as the solvent. When acids or bases are dissolved in water, they dissociate into their respective ions. For example:

• An acid like hydrochloric acid (HCl) will dissociate into hydrogen ions (H⁺) and chloride ions (Cl⁻).
• A base like sodium hydroxide (NaOH) will dissociate into sodium ions (Na⁺) and hydroxide ions (OH⁻).

This dissociation is crucial for the Arrhenius definition, as it highlights the increase in H⁺ or OH⁻ ions in the solution. The behavior of substances in an aqueous environment shows the fundamental aspect of acid-base chemistry. However, the specificity of this definition to aqueous solutions is also its major drawback, as it excludes reactions that occur outside of water. To fully understand the versatility of acid-base reactions, it's important to explore other definitions that are not confined to aqueous solutions.

Limitations of Arrhenius Definition

The Arrhenius definition of acids and bases, while useful, is quite limited. One significant limitation is its restriction to aqueous solutions. This definition fails to explain acid-base reactions that occur in non-aqueous solutions, such as organic solvents, or in gaseous states.
For example, chemical reactions involving ammonia (NH₃) and hydrogen chloride gas (HCl) do not adhere to the Arrhenius definition. In this case, NH₃ acts as a base by accepting a proton from HCl, which acts as an acid, forming ammonium chloride (NH₄Cl). No water is involved, and there are no H⁺ or OH⁻ ions produced in an aqueous solution.
Other definitions, like the Brønsted-Lowry and Lewis definitions, provide a broader understanding of acid-base behavior. The Brønsted-Lowry definition describes acids as proton donors and bases as proton acceptors, regardless of the solvent. The Lewis definition categorizes acids as electron pair acceptors and bases as electron pair donors, offering an even wider applicability. These definitions help explain many reactions that the Arrhenius definition cannot cover and are particularly important for advanced studies in chemistry.