In Class 11 Chemistry (Equilibrium chapter), the equilibrium constant can be expressed in terms of concentration (Kc) or partial pressure (Kp) for gaseous reactions. The mathematical relationship between them is frequently tested in exams.
When calculating Δn, you must strictly ignore solids and liquids. Only the moles of substances in the gaseous state (g) are counted.
The relationship is given by the equation:
Kp = Kc(RT)^Δn
Where:
The derivation relies on the Ideal Gas Equation (PV = nRT). Rearranging it gives P = (n/V)RT. Since (n/V) is moles per volume, it represents molar concentration (C). Therefore, P = CRT. By substituting this partial pressure into the Kp expression, we arrive at the final formula.
Kp is equal to Kc ONLY when Δn = 0. If the number of gaseous reactant moles equals the gaseous product moles, then Δn = 0. Since any number to the power of 0 is 1, the formula becomes Kp = Kc(1), so Kp = Kc. Example: H₂(g) + I₂(g) ⇌ 2HI(g). Here Δn = 2 - (1+1) = 0.
If Δn is positive (products have more gaseous moles than reactants), then **Kp > Kc**.
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