Question

Define Bronsted-Lowry base. How does it differ from an Arrhenius base?

Short answer

Bronsted-Lowry base accepts protons; Arrhenius base increases OH⁻ ions in water.

Step-by-step solution

Define Bronsted-Lowry Base

A Bronsted-Lowry base is defined as any substance that can accept a proton (H⁺ ion) during a chemical reaction. This definition relies on the presence of a proton donor from which the base can accept a proton.

Define Arrhenius Base

An Arrhenius base is defined as a substance that increases the concentration of hydroxide ions (OH⁻) in an aqueous solution. This definition is specific to aqueous solutions and does not consider other types of reactions.

Compare the Two Definitions

The primary difference between a Bronsted-Lowry base and an Arrhenius base lies in their definitions: a Bronsted-Lowry base focuses on the ability to accept protons, while an Arrhenius base releases hydroxide ions in water. Therefore, the Bronsted-Lowry theory is more general and can apply to reactions not occurring in solutions, unlike the Arrhenius definition.

Key concepts

Arrhenius Base

The Arrhenius base is a foundational concept in chemistry named after Svante Arrhenius. An Arrhenius base is characterized by its ability to increase hydroxide ions (OH⁻) in an aqueous solution. This means that when the base is dissolved in water, it dissociates or reacts to produce more OH⁻ ions, which are negatively charged.

• Common examples include sodium hydroxide (NaOH) and potassium hydroxide (KOH).
• They often feel slippery to the touch and alter the pH of solutions to become more basic or alkaline.
• In water, NaOH dissociates to Na⁺ and OH⁻, illustrating an Arrhenius base.

Arrhenius's theory is straightforward, but limited to reactions occurring in water. This means if a substance can increase hydroxide ion concentration, it fits the Arrhenius definition of a base. However, it does not encompass acid-base reactions outside of aqueous solutions.

Proton Acceptance

Proton acceptance is central to the Bronsted-Lowry theory, which encompasses broader reactions than just those occurring in water. In this theory, a base is any substance capable of accepting a proton (H⁺ ion). This isn't limited to solutions and can occur in any chemical reaction where proton transfer is possible.

• Ammonia (NH₃) is a classic example where it acts as a Bronsted-Lowry base by accepting a proton to become NH₄⁺ (ammonium).
• This concept applies to gas-phase reactions and other non-aqueous systems.
• It allows for a more comprehensive understanding of molecular interactions involving protons.

The flexibility of the Bronsted-Lowry theory helps explain a wider variety of chemical reactions, going beyond the scope of the Arrhenius definition. This generality is why the Bronsted-Lowry concept is more often used in advanced chemistry studies.

Hydroxide Ions

Hydroxide ions (OH⁻) are an integral part of understanding basic solutions, especially in the context of Arrhenius bases. These ions give bases their characteristic properties and are fundamental to pH calculations.

• OH⁻ ions result from the dissociation of bases in water, such as KOH or NaOH.
• They react with H⁺ ions to form water in neutralization reactions.
• The concentration of OH⁻ ions is directly related to the basicity of a solution.

Understanding the behavior and role of hydroxide ions helps in predicting the reactions of bases. Since they are a byproduct of base dissociation, their concentration affects whether a solution is acidic, neutral, or basic. If OH⁻ concentration exceeds that of H⁺ ions, the solution is basic. This fundamental property is pivotal to pH balance, titration, and other analytical techniques in chemistry.