← Labs|Chemical Equilibrium — Le Chatelier, ICE Table & Kc/Kp
Reaction
Initial Concentrations (M)
[N₂]₀1.00 M
[H₂]₀3.00 M
[NH₃]₀0.00 M
Le Chatelier Perturbations
N₂ + 3H₂ ⇌ 2NH₃
Industrial ammonia synthesis. Exothermic — lower T favours product but slows rate.
Kc
4.10e-4
Kp
6.86e-7
ΔH
-92 kJ
Δn(gas)
-2
Q vs Kc
Q < Kc
ICE Table (mol/L)
| N₂ | H₂ | NH₃ | |
|---|---|---|---|
| I (Initial) | +1.0000 | +3.0000 | +0.0000 |
| C (Change) | -0.0477 | -0.1431 | +0.0954 |
| E (Equilibrium) | +0.9523 | +2.8569 | +0.0954 |
Shift: Forward ⇒
|x| = 0.0477 M
Kc verified = 4.10e-4
Concentration vs Time (approach to equilibrium)
Equilibrium Laws
Kc Expression
For aA + bB ⇌ cC + dD:
Kc = [C]ᶜ[D]ᵈ / [A]ᵃ[B]ᵇ
Pure solids and liquids: concentration is constant → excluded from Kc expression
Kp = Kc × (RT)^Δn
Kp relates partial pressures
Δn = moles gas products − moles gas reactants
For this reaction: Δn = -2
R = 0.0821 L·atm/mol·K
Q vs Kc (Reaction Quotient)
Q < Kc → reaction goes FORWARD
Q > Kc → reaction goes REVERSE
Q = Kc → system AT equilibrium
Q uses any concentrations;
Kc uses equilibrium concentrations
Le Chatelier Principle
A system at equilibrium shifts to oppose any change:
• Add reactant → shift forward
• Add product → shift reverse
• ↑ T exo rxn → shift reverse
• ↑ T endo rxn → shift forward
• ↑ P: shift to fewer gas moles
• Catalyst: no shift, faster rate
ΔG and K
ΔG° = −RT ln K
K > 1: ΔG° < 0 (product-favoured)
K < 1: ΔG° > 0 (reactant-favoured)
At equilibrium: ΔG = 0
ΔG = ΔG° + RT ln Q
van't Hoff Equation
ln(K₂/K₁) = −ΔH/R × (1/T₂ − 1/T₁)
Exothermic: K↓ as T↑
Endothermic: K↑ as T↑
Plot ln K vs 1/T → slope = −ΔH/R