Differentiate Arrhenius and Bronsted-Lowry acid/base definitions and give an example of each.

• A. Arrhenius: acids and bases produce H+/OH− in water; Bronsted-Lowry: acids donate protons and bases accept protons (e.g., HCl; NH3).
• B. Arrhenius: acids donate protons and bases accept protons; Bronsted-Lowry: acids and bases are defined by proton transfer; e.g., NH3 as a base.
• C. Arrhenius: acids donate electrons; Bronsted-Lowry: acids donate protons.
• D. Arrhenius: focus on solvent effects; Bronsted-Lowry: focus on pH change.

Answer: (A)

The key idea is how acids and bases are defined. Arrhenius defines acids as substances that increase H+ in water (really H3O+ in solution) and bases as substances that increase OH− in water. Bronsted-Lowry defines acids as proton donors and bases as proton acceptors, and this proton-transfer view works regardless of the solvent.

This choice shows both ideas clearly. HCl in water donates a proton to the solvent, creating H3O+—an Arrhenius acid. It also donates a proton in Bronsted-Lowry terms, making it a Bronsted-Lowry acid. Ammonia accepts a proton to form NH4+, which is Bronsted-Lowry base behavior, even though NH3 does not produce OH− in water, so it isn’t an Arrhenius base. This illustrates why Bronsted-Lowry is more general: proton transfer defines acidity and basicity beyond aqueous solutions.