How to Memorize the Difference Between Arrhenius, Bronsted-Lowry, and Lewis Acids an Bases

Before memorizing the differences between the various definitions of acids and bases, take a look a closer look at the definitions themselves. Once familiar with them, you can move on to memorizing the specific distinctions.

The following will help you define and differentiate Arrhenius vs. Brønsted-Lowry vs. Lewis acids and bases.

Definitions of Acids and Bases

There are multiple definitions of acids and bases. The narrowest definition is the Arrhenius theory definition, which is primarily concerned with aqueous solutions.

An Arrhenius acid increases the concentration of H+ or H3O+ (hydronium) ions. Since protons don’t really float around in solution by themselves, hydronium is the more technically correct way to talk about protons in aqueous solution. An Arrhenius base increases the concentration of OH- ions.

An example of an Arrhenius acid is thus HCl. When HCl dissociates in solution, the hydronium ion concentration increases. An example of an Arrhenius base is NaOH. When NaOH dissociates in water, it increases the concentration of hydroxide ions.

By the Arrhenius definition: Acids release a proton, or H+, in water. Bases release a hydroxide ion, OH-, in water.

As stated previously, the Arrhenius theory definition of acids and bases is the narrowest since it only discusses aqueous solutions.

To be able to define more reactions, the Brønsted-Lowry definition focuses on proton transfer. A Brønsted-Lowry acid is any species that donates a proton to another molecule. A Brønsted-Lowry base is any species that accepts a proton from another molecule.

Finally, the Lewis definition is the broadest definition of acids and bases. Just as an Arrhenius acid is a Brønsted-Lowry acid, a Brønsted-Lowry acid is a Lewis acid.

In the Lewis definition, acids are electron pair acceptors. As a result of this, the acid is able to form a covalent bond with whatever supplies the electrons. Bases are electron pair donors.

Tips

    1. An Arrhenius acid increases the concentration of H+.
    2. An Arrhenius base increases the concentration of OH- ions.
    3. A Brønsted-Lowry acid is any species that donates a proton to another molecule. A Brønsted-Lowry base is any species that accepts a proton from another molecule.
    4. A Lewis acid is an electron pair acceptor. A Lewis base is an electron pair donor.

Tricks for Remembering the Difference

The great thing about the names of these definitions is that they are in alphabetical order going from most narrow to broadest definition. If you can keep in mind that:

Arrhenius < Brønsted-Lowry < Lewis

So, the first definition is the most narrow. Arrhenius only talks about aqueous solutions and whether or not a substance increases the hydronium or hydroxide ion concentration. Next is Brønsted-Lowry, which indicates that any substance that donates a proton is an acid, and anything that accepts it is a base. Finally, the Lewis definition is the broadest, stating that any electron pair acceptor is a Lewis acid, and an electron pair donor is a Lewis base.

Another trick is this: Arrhenius is all about the A’s. Arrhenius is concerned with AH ACID (a fun way of saying “an acid”). Here, the first A is Arrhenius and the H is a hydrogen or hydronium ion since the Arrhenius definition primarily concerns an increase in hydrogen ion concentration.

To recall the Lewis definition remember that the L is for Lewis and the E is for electrons (LEwis). The Lewis definition is primarily concerned with the movement of electrons.

Once you’ve got those two down, you know that the one that is left (Brønsted-Lowry definition) is concerned with the donation of protons.

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