Equilibria covers dynamic equilibrium, equilibrium constants (Kc, Kp), Le Chatelier\'s principle, and then acid-base chemistry — pH, Ka, buffers, titrations, and solubility product.
Dynamic equilibrium: rates of forward and reverse reactions are equal; concentrations are constant. Only in closed systems. Le Chatelier\'s principle: if conditions change, the equilibrium shifts to oppose the change. Temperature: increase favours endothermic direction (changes K). Pressure: increase favours fewer moles of gas. Concentration: increase shifts away from that side. Catalyst: no effect on position or K (speeds both equally). Equilibrium constant Kc: for aA + bB ⇌ cC + dD, Kc = [C]ᶜ[D]ᵈ / [A]ᵃ[B]ᵇ. Only changes with temperature. Kp: partial pressures, pₐ = xₐP (mole fraction × total pressure). Kp = (pC)ᶜ(pD)ᵈ / (pA)ᵃ(pB)ᵇ.
Brønsted-Lowry: acid = proton donor, base = proton acceptor. Strong acid: fully dissociates (HCl, HNO₃, H₂SO₄). Weak acid: partially dissociates (CH₃COOH). pH = −log[H⁺]. For strong acid: [H⁺] = concentration. For strong base: [OH⁻] = concentration, then Kw = [H⁺][OH⁻] = 1.0×10⁻¹⁴ at 25°C. For weak acid: Ka = [H⁺][A⁻]/[HA]; approximate [H⁺] = √(Ka × c). pKa = −log Ka. Buffer: resists pH change. Acidic buffer: weak acid + its conjugate base (e.g., CH₃COOH + CH₃COONa). Add acid → conjugate base neutralises; add base → weak acid neutralises. Henderson-Hasselbalch: pH = pKa + log([A⁻]/[HA]). Titration curves: strong-strong (sharp at pH 7), strong acid-weak base (endpoint < 7), weak acid-strong base (endpoint > 7). Indicator choice: pKa near equivalence pH. Ksp = [cation]ᵐ[anion]ⁿ for sparingly soluble salts.
An acidic buffer contains a weak acid (HA) and its conjugate base (A⁻, from a salt like NaA). The weak acid is mostly undissociated, providing a reservoir of HA. The salt fully dissociates, providing a reservoir of A⁻. When a small amount of acid (H⁺) is added: A⁻ + H⁺ → HA (the conjugate base mops up the H⁺, preventing pH drop). When a small amount of base (OH⁻) is added: HA + OH⁻ → A⁻ + H₂O (the weak acid neutralises the OH⁻, preventing pH rise). The pH remains approximately constant because both reservoirs are large. The Henderson-Hasselbalch equation pH = pKa + log([A⁻]/[HA]) shows pH is determined by the ratio — adding small amounts of acid or base barely changes this ratio. A buffer works best when [HA] ≈ [A⁻] (i.e., pH ≈ pKa).
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