Acids and Bases

Acids and Bases Definitions

  • Arrhenius Definition:

    • Acid: A substance that ionizes in water to produce hydronium ions (H₃O⁺).

    • Base: A substance that ionizes in water to produce hydroxide ions (OH⁻).

    • Example:

    • Strong Base: Ammonia (NH₃) acts as a base because it can accept a proton.

Brønsted-Lowry Concept

  • Proton Donor: An acid is defined as a substance that donates protons (ont[T]H⁺]ont!D*]).

  • Proton Acceptor: A base is defined as a substance that accepts protons.

  • Example Reactions:

    • NH₃(g) + H₂O(ℓ) ⇌ NH₄⁺(aq) + OH⁻(aq)

Strength of Acids and Bases

  • Strong Acids: Fully ionize in water, such as nitric acid (HNO₃).

  • Weak Acids: Only partially ionize in water, such as hydrofluoric acid (HF).

  • Conjugate acid-base pairs maintain their relationship through proton transfer.

  • Example: H₂O + NH₃ ⇌ NH₄⁺ + OH⁻ (Water acts as acid when reacting with NH₃)

Autoionization of Water

  • Reaction:

    • H₂O(ℓ) + H₂O(ℓ) ⇌ H₃O⁺(aq) + OH⁻(aq)

  • Ionization Constant:

    • Kᵂ = [H₃O⁺][OH⁻] = 1.0 x 10⁻¹⁴ (at 25°C)

Acid-Base Concentration Calculations

  • pH Calculation:

    • pH = −log₁₀[H₃O⁺]

    • Neutral solution at 25°C: pH = 7.00

    • Acidic solutions: pH < 7.00

    • Basic solutions: pH > 7.00

  • Example:

    • For a 6.0 M NaOH solution:

    • [H₃O⁺] = 1.0 x 10⁻¹⁴ / [OH⁻] = 1.7 x 10⁻¹⁵ M

Common Acid-Base Reactions

  • Neutral Reaction: A strong acid with a strong base yields a neutral solution e.g.

    • HCl + NaOH → NaCl + H₂O

  • Weak Acid with Strong Base:

    • Example: CH₃COOH + NaOH → CH₃COO⁻ + H₂O

    • pH of the resulting solution can be higher than 7 due to conjugate base formation.

Lewis Acids and Bases

  • Lewis Acid: Accepts a pair of electrons;

  • Lewis Base: Donates a pair of electrons.

  • Example: Metal ions can act as Lewis acids and form complex ions with Lewis bases.

Acid-Base Reactions in Daily Life

  • Antacids: Neutralize stomach acid (pH ≈ 1).

    • Example reactions include the reaction of Mg(OH)₂ with HCl.

  • Household Cleaners: Watch out for corrosive solutions with high/low pH.

Summary of Key Relationships

  • Stronger acids have weaker conjugate bases, and vice versa.

  • Primary acid strength depends on bond strength; weaker bonds lead to stronger acids.

  • Multiple acidic protons in polyprotic acids (e.g. H₃PO₄) exhibit different ionization constants (Ka) as each proton is lost.