Topic 7 covers reversible reactions and the factors that affect equilibrium position. SL uses Le Chatelier\'s principle qualitatively. HL introduces the equilibrium constant Kc and quantitative calculations.
In a closed system, a reversible reaction reaches equilibrium when the forward and reverse rates are equal. Macroscopic properties (concentrations, colour, pressure) remain constant. Equilibrium is dynamic — both reactions continue but at equal rates. Equilibrium can be approached from either direction.
If a system at equilibrium is disturbed, it shifts to partially counteract the disturbance. Increase concentration of reactant → shift right. Increase temperature: shift in endothermic direction. Increase pressure: shift towards fewer moles of gas. Catalyst: no effect on position — reaches equilibrium faster but same composition.
For aA + bB ⇌ cC + dD: Kc = [C]ᶜ[D]ᵈ / [A]ᵃ[B]ᵇ. Large Kc: products favoured. Small Kc: reactants favoured. Kc depends only on temperature. Pure solids and liquids are excluded from the expression. Reaction quotient Q has the same form but uses non-equilibrium concentrations. If Q < Kc, reaction proceeds forward; Q > Kc, proceeds in reverse.
For exothermic reactions: increasing T decreases Kc (shifts left). For endothermic reactions: increasing T increases Kc (shifts right). This is the only factor that changes Kc. Adding more reactant/product or changing pressure changes Q, not Kc — the system then shifts to restore Q = Kc.
No. A catalyst speeds up both the forward and reverse reactions equally, so equilibrium is reached faster but the equilibrium concentrations are unchanged. Kc is not affected. In industry, catalysts are important because they make equilibrium achievable in a practical timeframe — without a catalyst, a reaction might technically favour products but take too long to be useful.
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