14.2 Acid-Base Theories

Summary

This video explores the evolution of acid-base definitions, starting with the Arrhenius model, which focuses on aqueous solutions. It introduces the Bronsted-Lowry theory, which expands the concept to proton donors and acceptors, even outside water. The script also delves into Lewis acids and bases, broadening the scope even further to electron pair donors and acceptors. Examples are provided to illustrate how these models work with substances like ammonia, hydrochloric acid, and sulfuric acid. By examining various definitions, the video explains how acids and bases can fit multiple categories based on different chemical behaviors.

Takeaways

• 😀 The Arrhenius definition of acids and bases works in aqueous solutions, with acids releasing H⁺ ions and bases releasing OH⁻ ions.
• 😀 Bronsted-Lowry expanded the Arrhenius definition by defining acids as proton donors and bases as proton acceptors, applying beyond aqueous solutions.
• 😀 Some substances, like hydrochloric acid (HCl), act as acids outside of water, donating protons to substances like ammonia.
• 😀 Water itself can act as a Bronsted-Lowry acid when it donates a proton to ammonia, forming ammonium and hydroxide ions.
• 😀 The Bronsted-Lowry definition highlights the role of proton transfer, where acids donate protons and bases accept them.
• 😀 Sodium hydroxide (NaOH) is not strictly a Bronsted-Lowry base, but the hydroxide ion (OH⁻) it releases can act as one, accepting a proton from water.
• 😀 Monoprotic acids, like HCl, donate only one proton, while polyprotic acids, such as sulfuric acid (H₂SO₄), can donate multiple protons in a series of reactions.
• 😀 Sulfuric acid (H₂SO₄) ionizes in stages, first forming hydrogen sulfate (HSO₄⁻), which can then donate another proton to form sulfate (SO₄²⁻).
• 😀 Phosphoric acid (H₃PO₄) undergoes three stages of proton donation, eventually forming phosphate (PO₄³⁻).
• 😀 Lewis acids and bases expand the acid-base definition further, with acids accepting electron pairs and bases donating electron pairs, not necessarily involving protons.

Q & A

What was the original definition of acids and bases according to Arrhenius?

According to Arrhenius, acids were defined as substances that increase the concentration of H+ ions (protons) in an aqueous solution, while bases were substances that increase the concentration of OH- ions in an aqueous solution.

Why was the Arrhenius definition of acids and bases expanded?

The Arrhenius definition was expanded because some materials were found to behave like acids and bases outside of aqueous solutions, meaning they could donate or accept protons without water present.

How did the Bronsted-Lowry definition of acids and bases expand upon the Arrhenius definition?

The Bronsted-Lowry definition expanded the concept by defining acids as any molecule or ion that donates a proton (H+), and bases as any molecule or ion that accepts a proton, even outside of aqueous solutions.

What is the relationship between Arrhenius acids and Bronsted-Lowry acids?

All Arrhenius acids are also Bronsted-Lowry acids because they donate a proton (H+) in solution, but not all Bronsted-Lowry acids are Arrhenius acids, as the former can donate protons outside of aqueous solutions.

Can water act as a Bronsted-Lowry acid? If so, how?

Yes, water can act as a Bronsted-Lowry acid. When mixed with ammonia, water donates a proton (H+) to become a hydroxide ion (OH-) while leaving behind a hydronium ion (H3O+).

How does ammonia behave in a Bronsted-Lowry acid-base reaction?

Ammonia acts as a Bronsted-Lowry base because it accepts a proton (H+) from another substance, such as water, to form ammonium (NH4+).

What is the definition of a monoprotic acid, and can you provide an example?

A monoprotic acid is an acid that can donate only one proton (H+) per molecule. An example is hydrochloric acid (HCl), which donates only one proton in solution.

What is the difference between a monoprotic acid and a polyprotic acid?

A monoprotic acid donates only one proton, whereas a polyprotic acid can donate more than one proton. For example, sulfuric acid (H2SO4) is a polyprotic acid because it can donate two protons.

What is a Lewis acid, and how does it differ from a Bronsted-Lowry acid?

A Lewis acid is any atom or ion that accepts an electron pair to form a covalent bond. This differs from Bronsted-Lowry acids, which are defined by their ability to donate protons (H+).

Can a substance be both a Bronsted-Lowry acid and a Lewis acid? Provide an example.

Yes, a substance can be both a Bronsted-Lowry acid and a Lewis acid. For example, a proton (H+) is a Bronsted-Lowry acid because it donates a proton, and a Lewis acid because it accepts an electron pair to form a covalent bond with a base like ammonia.