Question

All Bronsted acids are Lewis acids, but the reverse is not true. Give two examples of Lewis acids that are not Bronsted acids.

Short answer

Boron trifluoride (BF3) and aluminum trichloride (AlCl3) are examples of Lewis acids that are not Bronsted acids. They are not Bronsted acids because they do not have a proton (H+) to donate.

Step-by-step solution

Identify Lewis Acids

First, let's understand Lewis acids. Lewis acids are electron pair acceptors. The examples are numerous, including simple substances like boron trifluoride (BF3), aluminum trichloride (AlCl3), and many metal cations like Ag+, Fe3+, and Sn4+.

Analyzing Lewis Acid as Bronsted Acid

Next, analyze whether these Lewis acids can donate a proton (H+) or not. For a compound to be a Bronsted acid, it must have a proton to donate. Neither BF3, AlCl3, Ag+, Fe3+, nor Sn4+ have a proton (H+) to donate.

Conclude the Results

From the analysis, it's clear that boron trifluoride (BF3) and aluminum trichloride (AlCl3) are examples of Lewis acids that are not Bronsted acids. They accept an electron pair but don't have a proton to donate.

Key concepts

Lewis Acid Examples

A Lewis acid is a chemical species that can accept an electron pair to form a new bond. In the given exercise, this definition helped identify substances like boron trifluoride (BF₃) and aluminum trichloride (AlCl₃) as Lewis acids because they can accept electron pairs through their empty orbitals.

Other examples of Lewis acids include metal cations such as silver ion (Ag⁺), iron(III) ion (Fe³⁺), and tin(IV) ion (Sn⁴⁺). These metal ions are able to accept electrons due to their positive charges and incomplete electron shells. It's important to note that Lewis acids are not restricted to compounds that have a positive charge; neutral molecules with central atoms that have vacant orbitals also act as Lewis acids.

Bronsted Acid Definition

A Bronsted acid, in contrast to a Lewis acid, is defined specifically as a substance that can donate a proton (H⁺) to another species. Proton donation is the hallmark of Bronsted acidity. This definition sets Bronsted acids apart because it focuses on the transfer of a proton in a chemical reaction, whereas Lewis acidity is defined by the ability to accept an electron pair, a broader criterion that is not limited to hydrogen-containing species. As such, all Bronsted acids are inherently Lewis acids, but not all Lewis acids are Bronsted acids, since the latter requires the presence of a transferable proton.

Proton Donation in Acids

In the realm of Bronsted-Lowry acid-base chemistry, proton donation is the primary characteristic of an acid. When a substance donates a proton, it behaves as a Bronsted acid. The proton is, essentially, a hydrogen atom that has lost its electron, leaving behind a positively charged hydrogen ion (H⁺).

This fundamental process can be seen in the behavior of common acids like hydrochloric acid (HCl), which readily donates a proton to a water molecule, forming a hydronium ion (H₃O⁺) and a chloride ion (Cl^-). Proton donation results in the formation of a conjugate base, which is what remains of the acid after the proton is donated.

Chemical Acid Classification

Chemical acids can be classified according to the definitions provided by Lewis and Bronsted. While Bronsted acids are characterized by their ability to donate a proton, Lewis acids are characterized by their ability to accept an electron pair. The classification allows chemists to predict and understand the behavior of substances in acid-base reactions.

Aside from Lewis and Bronsted acids, there are other categories based on strength (strong acids completely ionize in solution, whereas weak acids partially ionize), concentration (dilute or concentrated), and even by solubility (like hydrophilic and hydrophobic acids). To fully grasp chemical reactions, it's crucial for students to understand these classifications, as they guide the expected outcomes when acids interact with bases and other substances.