Chapter 14

Bronsted-Lowry Acids and Bases

• Definitions:

  • Acid: Donates a proton (H+)

  • Base: Accepts a proton (H+)

  • Conjugate pairs: Substances that differ only by the presence of a proton.

• Acid-Base Reaction:

  • Example: H2O + NH3 ⇌ NH4+ + OH−

Autoionization of Water

• Reaction: H2O ⇌ H3O+ + OH−

• Water can act as both an acid and a base (amphoteric).

• Ionization Constants:

  • Kw (Ion product of water) = [H3O+][OH−] = 1 x 10^-14

pH and pOH

• Calculations:

  • pH = -log[H3O+]

  • pOH = -log[OH−]

• Solution Types:

  • Acidic: [H3O+] > [OH−]

  • Basic: [OH−] > [H3O+]

  • Neutral: [H3O+] = [OH−]

Relative Strength of Acids and Bases

• Strong Acids/Bases: Completely ionize e.g. HCl + H2O ⇌ H3O+ + Cl−, NaOH → Na+ + OH−

• Weak Acids/Bases: Partially ionize, strength indicated by Ka for acids and Kb for bases.

  • Example: HA + H2O ⇌ H3O+ + A− (Ka = [H3O+][A−]/[HA])

Trends in Acid Strength

• Factors affecting acidity:

  • Bond strength: Weaker bonds indicate stronger acids.

  • Polarity of the bond: More polar indicates stronger acids.

  • Oxyacids: Stronger with more electronegative central atoms and higher oxidation states.

Acid-Base Equilibrium Calculations

• Steps:

  1. Find Ka or Kb from equilibrium concentrations.

  2. Find Ka or Kb from pH and initial concentrations.

  3. Calculate equilibrium concentrations and pH from Ka or Kb.

• Example Calculation:

  • Given: HNO2 + H2O ⇌ H3O+ + NO2−, with a calculated pH leading to [H3O+] = 0.0046 M, Ka = 0.0470.