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This lesson covers the equilibrium constant Kc, writing Kc expressions, ICE tables, calculating Kc, its units, and the effect of changing conditions on Kc.
For a general reaction: aA + bB ⇌ cC + dD
The equilibrium constant in terms of concentrations is:
Kc = [C]^c [D]^d / ([A]^a [B]^b)
where [X] represents the equilibrium concentration of species X in mol dm⁻³.
Example 1: N₂(g) + 3H₂(g) ⇌ 2NH₃(g)
Kc = [NH₃]² / ([N₂][H₂]³)
Example 2: CH₃COOH(l) + C₂H₅OH(l) ⇌ CH₃COOC₂H₅(l) + H₂O(l)
Kc = [CH₃COOC₂H₅][H₂O] / ([CH₃COOH][C₂H₅OH])
Note: Although these are liquids, they are in a homogeneous mixture (not pure liquids), so all species are included.
An ICE table (Initial, Change, Equilibrium) is used to track concentrations through a reaction reaching equilibrium.
Example 3: 1.0 mol of N₂O₄ is placed in a 1.0 dm³ container at 25°C. At equilibrium, 0.20 mol of N₂O₄ has dissociated. Calculate Kc for: N₂O₄(g) ⇌ 2NO₂(g)
| N₂O₄ | NO₂ | |
|---|---|---|
| Initial (mol) | 1.0 | 0 |
| Change (mol) | −0.20 | +0.40 |
| Equilibrium (mol) | 0.80 | 0.40 |
Since volume = 1.0 dm³: [N₂O₄] = 0.80 mol dm⁻³, [NO₂] = 0.40 mol dm⁻³
Kc = [NO₂]² / [N₂O₄] = (0.40)² / 0.80 = 0.16 / 0.80 = 0.20 mol dm⁻³
The units of Kc depend on the balanced equation. Substitute the units (mol dm⁻³) for each concentration term.
For Kc = [NO₂]² / [N₂O₄]:
Units = (mol dm⁻³)² / (mol dm⁻³) = mol dm⁻³
For Kc = [NH₃]² / ([N₂][H₂]³):
Units = (mol dm⁻³)² / ((mol dm⁻³)(mol dm⁻³)³) = (mol dm⁻³)² / (mol dm⁻³)⁴ = mol⁻² dm⁶
Exam Tip: Always work out the units of Kc. If the total powers of concentration on top and bottom are equal, Kc has no units.
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