Lecture Notes on Acids and Bases

Introduction

• This module focuses on acids and bases as an application of equilibrium.
• Concepts from acid-base chemistry in lab contexts will be revisited.
• Emphasis on how equilibrium principles apply to acids and bases.

Definitions of Acids and Bases

Arrhenius Definition

• Acids: Yield hydronium ions (H3O+) in aqueous solution.
• Bases: Yield hydroxide ions (OH-) in aqueous solution.

Bronsted-Lowry Definition

• Acids: Donate a hydrogen ion (H+) to another species.
• Bases: Accept a hydrogen ion (H+).

Lewis Definition

• Will be explored later on.

General Characteristics

• Acids: pH < 7 in aqueous solution.
• Bases: pH > 7 in aqueous solution.
• Neutral Solution: pH = 7 with equal concentrations of H+ and OH-.

Reactions and Examples

• Sodium Hydroxide (NaOH)
  • Splits into Na+ and OH- in solution: classified as a base (Arrhenius).
• Hydrochloric Acid (HCl)
  • Reacts with water to form H3O+ and Cl-: classified as an acid (Arrhenius & Bronsted-Lowry).

Water as Solvent

• All reactions considered involve water as the solvent (aqueous systems).
• Different rules may apply in non-aqueous systems.

Amphiprotic Nature of Water

• Amphiprotic/Amphoteric: Can act as both an acid and a base.
  • Example: Water can donate or accept a proton depending on the reaction context.

Conjugate Acid-Base Pairs

• NH3 and NH4+
  • NH3 is the base; NH4+ is the conjugate acid.
• H2O and OH-
  • H2O is the acid; OH- is the conjugate base.

Identifying Conjugate Pairs

Example Reactions

• H2SO4 + H2O → HSO4- + H3O+

  • H2SO4 is the acid; HSO4- is the conjugate base.
  • H2O is the base; H3O+ is the conjugate acid.

• HCO3- + H2O ⇌ H2CO3 + OH-

  • HCO3- is the base; H2CO3 is the conjugate acid.
  • H2O is the acid; OH- is the conjugate base.

Conclusion

• Focus on understanding how acids and bases relate to equilibrium.
• Importance of identifying the role of substances in reactions (acid or base).
• Conjugate pairs help connect acids and bases within reaction frameworks.