Lecture on Acids and Bases

Definitions

Arrhenius Definition

• Acids: Release H+ ions (hydronium ions, H3O+) in solution.
• Bases: Release hydroxide ions (OH-) in solution.
• Examples:
  • Arrhenius Acids: HF, HCl, H2SO4, HNO3
  • Arrhenius Bases: NaOH, KOH, Ca(OH)2

Bronsted-Lowry Definition

• Acids: Proton donors (release H+ ions).
• Bases: Proton acceptors.
• Reaction Example: HF (acid) + H2O (base) ↔ H3O+ (conjugate acid) + F- (conjugate base)
  • Conjugate Acid: Formed by adding a hydrogen ion.
  • Conjugate Base: Formed by removing a hydrogen ion.

Conjugate Acid-Base Pair Examples

• Bicarbonate Ion (HCO3-)
  • Conjugate Acid: H2CO3
  • Conjugate Base: CO3 2-
• Ammonia (NH3)
  • Conjugate Acid: NH4+
  • Conjugate Base: NH2-
• Hydrogen Phosphate (HPO4 2-)
  • Conjugate Acid: H2PO4-
  • Conjugate Base: PO4 3-

Reaction Examples (identifying Bronsted-Lowry Acids/Bases)

• NH3 + H2O ↔ NH4+ + OH-
  • NH3: Base (proton acceptor)
  • H2O: Acid (proton donor)
  • NH4+: Conjugate Acid
  • OH-: Conjugate Base
• CH3OH + H2O ↔ CH3O- + H3O+
  • CH3OH: Acid (proton donor)
  • H2O: Base (proton acceptor)
  • CH3O-: Conjugate Base
  • H3O+: Conjugate Acid

Lewis Acid-Base Definition

• Lewis Acid: Electron pair acceptor.
• Lewis Base: Electron pair donor.
• Example: Reaction between BH3 and NH3.
  • Boron (BH3): Lewis acid (electron poor, accepts electron pair).
  • Nitrogen (NH3): Lewis base (electron rich, donates electron pair).

Additional Concepts

• Lewis Base: Nucleophile (electron rich).
• Lewis Acid: Electrophile (electron poor).

These definitions help us understand the behavior of different substances in chemical reactions and their roles in forming new compounds.