Question

If a substance is a Lewis acid, is it necessarily a BronstedLowry acid? Is it necessarily an Arrhenius acid?

Short answer

A substance being a Lewis acid does not necessarily mean it is a Bronsted-Lowry acid or an Arrhenius acid. While some Lewis acids can also be Bronsted-Lowry or Arrhenius acids, the definitions and requirements for each type of acid are distinct. Lewis acids accept electron pairs, Bronsted-Lowry acids donate protons, and Arrhenius acids dissociate in water to produce hydrogen ions.

Step-by-step solution

Define Lewis Acids

Lewis acids are substances that can accept electron pairs, typically forming a covalent bond with a Lewis base, which donates the electron pair. These substances are not limited to hydrogen-containing molecules, unlike Bronsted-Lowry and Arrhenius acids.

Define Bronsted-Lowry Acids

A Bronsted-Lowry acid is a substance that donates a proton (hydrogen ion, H+) to another species (called a Bronsted-Lowry base) in a chemical reaction. This definition is broader than that of an Arrhenius acid because it doesn't require the presence of a solvent.

Define Arrhenius Acids

Arrhenius acids are substances that dissociate in water to produce hydrogen ions (H+), which combine with water molecules to form hydronium ions (H3O+). This definition limits itself to reactions that occur in aqueous solutions.

Compare Lewis acids to Bronsted-Lowry acids

A Lewis acid can be a Bronsted-Lowry acid if it accepts an electron pair and also donates a proton (H+). However, not all Lewis acids are Bronsted-Lowry acids, as some might not have a hydrogen atom to donate as a proton. Therefore, a substance being a Lewis acid does not necessarily mean it is a Bronsted-Lowry acid.

Compare Lewis acids to Arrhenius acids

Similarly, a Lewis acid can be an Arrhenius acid if it accepts an electron pair and dissociates in water to produce hydrogen ions (H+). However, not all Lewis acids are Arrhenius acids because some might not produce hydrogen ions in aqueous solution. Therefore, a substance being a Lewis acid does not necessarily mean it is an Arrhenius acid. In conclusion, being a Lewis acid does not necessarily mean that the substance is also a Bronsted-Lowry acid or an Arrhenius acid. While there may be some overlap, the definitions and requirements for each type of acid are distinct.

Key concepts

Lewis acids

In acid-base chemistry, Lewis acids stand out by their unique capability to accept electron pairs. This property allows them to form covalent bonds with Lewis bases, which are the electron pair donors. Unlike more conventional definitions involving protons, Lewis acids don't necessarily include hydrogen atoms.
For example, aluminum chloride (AlCl₃) or boron trifluoride (BF₃) are classified as Lewis acids because they can accept an electron pair despite lacking hydrogen. This means they can act as acids in reactions without the need for a solvent or to donate protons. This broader definition suits a wide range of chemical reactions, making Lewis acids predominant in non-aqueous media.

Bronsted-Lowry theory

According to the Bronsted-Lowry theory, acids are defined by their ability to donate protons (H⁺ ions). This theory expanded the concept of acids beyond the Arrhenius definition by focusing on proton donors and acceptors, rather than requiring the acid to increase H⁺ concentration in water.
This allows for a more universal approach to identifying acids and bases, as the need for the presence of water as a solvent is no longer a restriction. Bronsted-Lowry acids must have a hydrogen atom to donate, which limits this category in comparison to Lewis acids. While not every Lewis acid can be a Bronsted-Lowry acid, every Bronsted-Lowry acid is inherently part of Lewis's broader framework since donating a proton involves electron pair acceptance.

Arrhenius theory

Arrhenius theory defines acids in terms of their behavior in aqueous solutions. Specifically, an Arrhenius acid is a substance that dissociates to release hydrogen ions (H⁺) into the water, leading to the formation of hydronium ions (H₃O⁺). This definition is restricted to reactions that happen in water.
The need for aqueous solutions in the Arrhenius definition limits it compared to Bronsted-Lowry and Lewis theories. Despite these limitations, Arrhenius acid-base reactions are still fundamental to understanding many chemical processes that occur in water, such as those in biological systems and industrial reactions. Arrhenius acids are inherently Bronsted-Lowry acids since dissociating in water involves proton donation.