By Riti Gupta, Updated Mar 24, 2022

Before committing the differences among the three main acid–base frameworks to memory, first understand each definition’s core concepts. Once you’re comfortable with the fundamentals, the distinctions become much easier to recall.

Core Definitions of Acids and Bases

There are several ways to define acids and bases, ranging from the most specific to the most inclusive. Below is a concise comparison of the Arrhenius, Brønsted–Lowry, and Lewis frameworks.

Arrhenius

• Applies only to aqueous solutions.
• Acid: Substance that increases the concentration of α⁻³ (hydronium) ions.
• Base: Substance that increases the concentration of δ⁻³ (hydroxide) ions.
• Example acid: HCl (hydrochloric acid).
• Example base: NaOH (sodium hydroxide).

Brønsted–Lowry

• Broader than Arrhenius; focuses on proton transfer.
• Acid: Any species that donates a proton to another molecule.
• Base: Any species that accepts a proton from another molecule.

Lewis

• Most inclusive; based on electron pair interactions.
• Acid: Electron‑pair acceptor, capable of forming a covalent bond with an electron donor.
• Base: Electron‑pair donor.

TL;DR (Quick Take)

• Arrhenius: Acid increases H⁻³; base increases OH⁻³.
• Brønsted–Lowry: Acid donates a proton; base accepts a proton.
• Lewis: Acid accepts an electron pair; base donates an electron pair.

Mnemonic Tricks to Keep the Definitions Straight

The order of the terms—Arrhenius, Brønsted–Lowry, Lewis—mirrors their breadth: from narrow to broad. Remember:

• Arrhenius less than Brønsted–Lowry less than Lewis.
• Arrhenius: Think Acid, Hydrogen—focus on hydronium concentration.
• Brønsted–Lowry: Centered on proton donation/acceptance.
• Lewis: Leans on Electrons—electron pair transfer.

By anchoring each definition to a single key concept—hydronium for Arrhenius, proton transfer for Brønsted–Lowry, and electron pairs for Lewis—you’ll be able to recall the differences instantly.