Question

Give Arrhenius's and Bronsted's definitions of an acid and a base. Why are Bronsted's definitions more useful in describing acid-base properties?

Short answer

Bronsted's definitions are more useful because they apply to more environments and aren't limited to aqueous solutions.

Step-by-step solution

Understanding Arrhenius's Definitions

Arrhenius defines an acid as a substance that increases the concentration of hydrogen ions (\( H^+ \)) in an aqueous solution. Conversely, an Arrhenius base is defined as a substance that increases the concentration of hydroxide ions (\( OH^- \)) in an aqueous solution.

Understanding Bronsted's Definitions

Bronsted defines an acid as a proton (\( H^+ \)) donor and a base as a proton (\( H^+ \)) acceptor. This definition is not limited to aqueous solutions and is based on the transfer of protons between substances.

Comparing the Two Definitions

The Bronsted definitions are more universal and applicable because they are not restricted to aqueous solutions. They extend the concept of acids and bases to reactions that involve solvents other than water and allow the description of acid-base behavior in non-aqueous environments.

Key concepts

Arrhenius Acid

In acid-base chemistry, understanding the fundamentals starts with Arrhenius acids. Arrhenius defined an acid as a substance that, when dissolved in water, increases the concentration of hydrogen ions, \( H^+ \). Simply put, when an Arrhenius acid is in a solution, it releases hydrogen ions. For example, hydrochloric acid \( (\text{HCl}) \) dissociates in water to produce hydrogen ions and chloride ions.

• Key Point 1: Arrhenius acids are identified by their ability to increase hydrogen ion concentration in water.
• Key Point 2: This definition is limited to aqueous (water-based) environments.

While the Arrhenius concept is foundational, it's specific to scenarios where water is the solvent, limiting its application to certain reactions.

Arrhenius Base

Moving on to Arrhenius bases, these are substances that increase the concentration of hydroxide ions, \( OH^- \), when dissolved in water. For instance, sodium hydroxide \( (\text{NaOH}) \) quand disolved, it dissociates into sodium and hydroxide ions, leading to a greater concentration of hydroxide ions in the solution.

• Key Point 1: Arrhenius bases increase hydroxide ion concentration in aqueous solutions.
• Key Point 2: The definition limits the functionality of Arrhenius bases to only water-based solutions.

Although beneficial for simple reactions, the Arrhenius base definition does not accommodate reactions in non-water solvents.

Bronsted Acid

A Bronsted acid provides a broader perspective by being a proton \( H^+ \) donor. Rather than being confined to water, a Bronsted acid can donate protons to any base. Consider hydrochloric acid again, in the presence of a base, the acid donates its proton to the base.

• Key Point 1: Bronsted acids are defined by their ability to donate protons.
• Key Point 2: This concept is not limited to aqueous solutions but is applicable in varied solvent environments.

The flexibility of the Bronsted definition allows it to describe more chemical reactions beyond those possible in simple aqueous settings.

Bronsted Base

A Bronsted base complements its acid counterpart by acting as a proton acceptor. This means that a Bronsted base gains a proton from an acid during a reaction. Ammonia \( (\text{NH}_3) \), for example, when reacting with water, accepts a proton and becomes ammonium \( (\text{NH}_4^+) \).

• Key Point 1: Bronsted bases accept protons from acids, broadening the definition of bases beyond just hydroxide producers.
• Key Point 2: Because it is not limited to water, the Bronsted base supports a wide range of chemical environments, including those in organic solvents.

The flexibility of the Bronsted approach is key for describing and predicting the behavior of substances in various chemical reactions.