Units 7-8 cover chemical equilibrium — when forward and reverse rates are equal — and the major application: acid-base chemistry including pH calculations, buffers, and titrations.
At equilibrium: rate forward = rate reverse, concentrations constant. K expression: products over reactants, each raised to stoichiometric power. Pure solids and liquids excluded. Q (reaction quotient): same form as K but with current concentrations. Q < K: reaction proceeds forward. Q > K: proceeds in reverse. Q = K: at equilibrium. ICE tables: Initial, Change, Equilibrium — systematic method for calculations. K is temperature-dependent only.
System at equilibrium responds to oppose change. Add reactant → shifts right (makes more product). Remove product → shifts right. Increase pressure → shifts to side with fewer gas moles. Temperature: increase T → shifts in endothermic direction AND changes K value. Catalyst: no effect on position or K (reaches equilibrium faster).
Brønsted-Lowry: acid = H⁺ donor, base = H⁺ acceptor. Strong acids/bases: fully dissociate. Weak: partially dissociate — use Ka/Kb. pH = -log[H⁺]. Kw = [H⁺][OH⁻] = 1.0×10⁻¹⁴ at 25°C. Ka × Kb = Kw for conjugate pairs. Buffer: weak acid + conjugate base. Resists pH change. Henderson-Hasselbalch: pH = pKa + log([A⁻]/[HA]). Titration curves: equivalence point where moles acid = moles base. Half-equivalence: pH = pKa.
A buffer contains a weak acid (HA) and its conjugate base (A⁻) in significant concentrations. When H⁺ is added: A⁻ + H⁺ → HA (the conjugate base neutralises the acid). When OH⁻ is added: HA + OH⁻ → A⁻ + H₂O (the weak acid neutralises the base). The ratio [A⁻]/[HA] changes only slightly, so pH changes minimally (Henderson-Hasselbalch). Buffer capacity: the more moles of HA and A⁻ present, the more acid/base the buffer can absorb. Buffer is most effective when pH ≈ pKa (equal amounts of HA and A⁻).
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