Acid-Base Equilibria Notes

Chemistry: The Central Science - Chapter 16: Acid-Base Equilibria

Definitions of Acids and Bases

  • Arrhenius Definition:

  • Acid: Substance that increases H⁺ concentration in water.

  • Base: Substance that increases OH⁻ concentration in water.

  • Brønsted–Lowry Definition:

  • Acid: Proton (H⁺) donor.

  • Base: Proton (H⁺) acceptor.

  • Lewis Definition:

  • Lewis Acid: Electron pair acceptor.

  • Lewis Base: Electron pair donor.

Water as a Proton Acceptor

  • Formation of hydronium ion (H₃O⁺) when H⁺ is dissolved in water, which is often represented as [H⁺] = [H₃O⁺].

Brønsted–Lowry Acids and Bases

  • Brønsted–Lowry acids must have an available proton (H⁺) to donate.
  • Brønsted–Lowry bases must have at least one lone pair of electrons to accept a proton.

Amphiprotic Nature of Water

  • Water can behave both as an acid (donates H⁺) and a base (accepts H⁺):
  • Example: HCl + H2O → Cl⁻ + H3O⁺ (acting as base)
  • Example: NH₃ + H2O ⇌ NH₄⁺ + OH⁻ (acting as acid).

Conjugate Acids and Bases

  • Conjugate Acid-Base Pair: Two species that differ by one H⁺.
  • Removing H⁺ from an acid yields its conjugate base.
  • Adding H⁺ to a base yields its conjugate acid.
  • Example: from HClO₄ (acid) → ClO₄⁻ (base); from CN⁻ (base) → HCN (acid).

Relative Strengths of Acids and Bases

  • Strong Acids: Acids above H₃O⁺ are strong acids (e.g., HCl, HNO₃).
  • Strong Bases: Bases below OH⁻ are strong bases (e.g., NaOH, KOH).
  • Acids and bases that are weaker will yield their conjugate pairs.

Equilibrium of Acid-Base Reactions

  • Equilibrium favors the side with the weaker acid and base.
  • Example reactions show how to predict the side favored by equilibrium.

Autoionization of Water

  • Water can self-ionize, revealing its amphiprotic nature:
  • H₂O(l) + H₂O(l) ⇌ H₃O⁺(aq) + OH⁻(aq).

Ion Product Constant (Kₕ)

  • For pure water at 25°C: K_w = [H⁺][OH⁻] = 1.0 × 10⁻¹⁴.

Acidic, Basic, or Neutral Solutions

  • Neutral solution: [H⁺] = [OH⁻]
  • Acidic solution: [H⁺] > [OH⁻]
  • Basic solution: [H⁺] < [OH⁻].

pH Scale

  • pH: Measurement of H⁺ concentration.
  • pH = -log[H⁺]
  • Neutral pH = 7, Acidic pH < 7, Basic pH > 7.

Calculating pH

  • Example: [OH⁻] = 1.8 × 10⁻⁹ M; calculate pH using Kₕ relationship: pH = 14 - pOH.

Strong Acids and Bases

  • Strong acids (e.g., HCl, H₂SO₄) dissociate completely in solution.
  • Strong bases (e.g., NaOH, KOH) also dissociate completely.

Weak Acids and Bases

  • Weak acids do not fully dissociate.
  • Ka: Acid dissociation constant indicates strength of weak acids.
  • Weak bases have a similar constant called Kb.

Polyprotic Acids

  • Can donate more than one proton; each dissociation can be analyzed with K₁, K₂, etc.

Acid-Base Properties of Salts

  • Determine acidity or basicity based on cation and anion properties:
  1. Cations of strong bases are neutral.
  2. Small, highly charged metals are acidic.
  3. Anions of strong acids are neutral, while weak acids’ conjugates are basic.

Lewis Acid-Base Chemistry

  • Lewis acids accept electrons; Lewis bases donate electrons.
  • Examples include reactions involving hydrated metal cations.