Revise: Chemical Equilibrium
What dynamic equilibrium means, how Le Chatelier's principle predicts shifts from concentration, pressure, and temperature changes, and how the equilibrium constant Kc quantifies a reaction's balance point.
At equilibrium, forward rate = reverse rate — both reactions keep happening.
A saturated sugar solution: dissolving and crystallizing continue at equal rates.
A disturbed equilibrium shifts to partially counteract the disturbance.
Adding N₂ to N₂+3H₂⇌2NH₃ shifts the equilibrium forward.
Higher pressure favours the side with fewer gas moles.
N₂+3H₂⇌2NH₃: higher pressure favours NH₃ (2 moles vs. 4).
Kc = [products]^coef / [reactants]^coef, fixed at a given temperature.
[HI]=1.60, [H₂]=[I₂]=0.20 → Kc = 64.
Kc >> 1 favours products; Kc << 1 favours reactants.
Kc = 10¹⁵ → reaction nearly complete.
Reversible reactions run both forward and backward, shown with ⇌.
CO₂ + H₂O ⇌ H₂CO₃ in sealed vs. opened soda.
Equilibrium ≠ equal concentrations — it means equal RATES.
A saturated solution keeps dissolving and crystallizing at equal rates.
The system shifts to PARTIALLY offset a disturbance, reaching a new equilibrium.
Haemoglobin binds O₂ in lungs, releases it in tissues.
Adding a substance shifts equilibrium away from it; removing shifts toward it.
Removing NH₃ as it forms shifts N₂+3H₂⇌2NH₃ forward.
Higher temperature shifts equilibrium in the endothermic direction.
Raising temperature on an exothermic forward reaction favours the reverse.
Coefficients in the balanced equation become exponents in Kc.
[HI]=0.80, [H₂]=[I₂]=0.10 → Kc = 64.
Kc ≈ 1 means significant amounts of both reactants and products at equilibrium.
Kc = 10⁻¹² → reaction barely proceeds at all.