Chemistry UNIT 3 Topic 4,5,6

Properties of acids and bases

Acids pH < 7 donate H⁺; Bases pH > 7 accept H⁺; Strong acids/bases fully dissociate → high conductivity; Weak acids/bases partially dissociate → low conductivity; Strong/weak = dissociation; Concentrated/dilute = amount of solute

Strong acids

HCl, HNO₃, H₂SO₄

Weak acids

CH₃COOH, H₂CO₃

Strong bases

NaOH, KOH, Ba(OH)₂

Weak bases

NH₃, amines (R–NH₂)

Monoprotic, diprotic, triprotic

Monoprotic donates 1 H⁺; Diprotic donates 2 H⁺; Triprotic donates 3 H⁺

Self‑ionisation of water

2H₂O ⇌ H₃O⁺ + OH⁻; Kw = 1×10⁻¹⁴ at 25°C; Neutral water [H⁺] = [OH⁻] = 1×10⁻⁷ M

pH and pOH relationships

pH = –log[H⁺]; pOH = –log[OH⁻]; pH + pOH = 14

Brønsted–Lowry acid and base

Acid donates H⁺; Base accepts H⁺

Amphiprotic substances

Can donate or accept H⁺; Examples

Conjugate acid–base pairs

Acid → conjugate base (minus H⁺); Base → conjugate acid (plus H⁺); Example

Buffer definition

Weak acid + conjugate base OR weak base + conjugate acid; resists pH change

How buffers resist pH change

Added acid → buffer absorbs H⁺; Added base → buffer donates H⁺

Ka and Kb formulas

Ka = [H₃O⁺][A⁻] / [HA]; Kb = [BH⁺][OH⁻] / [B]

pKa and pKb

pKa = –log Ka; pKb = –log Kb; Lower pKa = stronger acid; Lower pKb = stronger base

Kw relationship

Kw = Ka × Kb

Weak acid/base pH calculation

Use ICE table; Use Ka or Kb; Assume x ≪ initial concentration; Solve for [H⁺] or [OH⁻]