Acids and Bases (Bronsted-Lowry Acids and Bases)

Introduction to Acids and Bases

Acids and bases are fundamental chemical species with distinct properties and behaviors in aqueous solutions. Their definitions have evolved to include a broader range of substances, especially with the Bronsted-Lowry theory.

What are Acids?

Definition and Properties

• Acids are substances that produce hydrogen ions (H+) when dissolved in water.

• They typically have a sour taste (e.g., vinegar, lemon juice).

• Acids turn blue litmus paper red and can corrode some metals.

• They are characterized by their ability to donate protons (H+ ions).

Naming Acids

• Simple Anions: If the acid contains a simple nonmetal anion, use the prefix "hydro-" and the suffix "-ic acid." Example: HCl (aq) → hydrochloric acid

• Polyatomic Ions:

  • If the polyatomic ion ends in "-ate," use "-ic acid." Example: HNO3 (nitrate) → nitric acid

  • If the polyatomic ion ends in "-ite," use "-ous acid." Example: HNO2 (nitrite) → nitrous acid

What is a Base?

Definition and Properties

• Bases are ionic compounds that dissociate into cations and hydroxide ions (OH-) in water.

• Most bases are formed from Group 1 and Group 2 metals (e.g., NaOH, KOH, Ca(OH)2).

• Bases have a bitter taste and a slippery feel.

• Bases turn litmus paper blue and phenolphthalein indicator pink.

Naming Bases

• Name the metal first, followed by "hydroxide." Examples: LiOH → lithium hydroxide, NaOH → sodium hydroxide

Bronsted-Lowry Acids and Bases

Expanded Definitions

The Bronsted-Lowry theory (1923) defines acids and bases based on proton transfer, not just the presence of hydroxide ions.

• Bronsted-Lowry Acid: Proton (H+) donor

• Bronsted-Lowry Base: Proton (H+) acceptor

In aqueous solutions, free hydrogen ions do not exist independently; they associate with water molecules to form the hydronium ion (H3O+).

Examples of Bronsted-Lowry Acid-Base Reactions

• HCl + H2O → H3O+ + Cl- - HCl is the acid (proton donor), H2O is the base (proton acceptor).

• NH3 + H2O ⇌ NH4+ + OH- - H2O is the acid (proton donor), NH3 is the base (proton acceptor).

Conjugate Acid-Base Pairs

Definition and Identification

In every acid-base reaction, two conjugate acid-base pairs are present. These pairs differ by one proton (H+).

• The acid loses a proton to become its conjugate base.

• The base gains a proton to become its conjugate acid.

Example: In the reaction HF + H2O ⇌ F- + H3O+:

• HF (acid) / F- (conjugate base)

• H2O (base) / H3O+ (conjugate acid)

Strength of Acids and Bases

Strong vs. Weak Acids

• Strong acids dissociate completely in water, producing a high concentration of H3O+ ions.

• Weak acids only partially dissociate, producing fewer H3O+ ions.

Strong vs. Weak Bases

• Strong bases dissociate completely in water, yielding a high concentration of OH- ions.

• Weak bases are poor acceptors of H+ and produce few ions in solution.

Summary Table: Properties of Acids and Bases

Key Equations

• Acid dissociation:

• Base dissociation:

References

• Timberlake, K. (2018). Chemistry: Introduction to General, Organic, and Biological Chemistry (13th ed). Pearson Education.