Equilibrium Constant
The equilibrium constant (K) is a dimensionless number that expresses the ratio of the concentrations (or partial pressures) of products to reactants, each raised to the power of their stoichiometric coefficients, for a reversible reaction at equilibrium at a given temperature. A large K (K >> 1) indicates the equilibrium favours products, while a small K (K << 1) indicates reactants predominate. K changes with temperature but is independent of initial concentrations, catalysts, or pressure (for Kc).
Key Formula
Kc = ([C]^c × [D]^d) / ([A]^a × [B]^b)
LaTeX: K_c = \frac{[C]^c[D]^d}{[A]^a[B]^b}
| Symbol | Meaning | Unit |
|---|---|---|
| [C], [D] | Molar concentrations of products C and D at equilibrium | mol/L |
| [A], [B] | Molar concentrations of reactants A and B at equilibrium | mol/L |
| a, b, c, d | Stoichiometric coefficients from the balanced equation | dimensionless |
Worked Example
Problem
For the reaction N₂(g) + 3H₂(g) ⇌ 2NH₃(g), at equilibrium at 500 °C the concentrations are: [N₂] = 0.50 mol/L, [H₂] = 0.15 mol/L, [NH₃] = 0.25 mol/L. Calculate Kc.
Solution
Step 1 – Write the Kc expression: Kc = [NH₃]² ÷ ([N₂][H₂]³). Step 2 – Substitute values: Kc = (0.25)² ÷ (0.50 × (0.15)³). Step 3 – Numerator: 0.25² = 0.0625. Step 4 – Denominator: 0.50 × 0.003375 = 0.0016875. Step 5 – Kc = 0.0625 ÷ 0.0016875 = 37.0.
Answer
Kc = 37.0 (dimensionless)
Interpreting the Magnitude of K
| Value of K | Equilibrium Position | Reaction Tendency | Example |
|---|---|---|---|
| K >> 10³ | Far right (products) | Reaction nearly complete | Strong acid dissociation |
| K = 1–10³ | Favours products | Products predominate | Many industrial reactions |
| K ≈ 1 | Middle | Comparable amounts of both | N₂O₄ ⇌ 2NO₂ at ~55 °C |
| K = 10⁻³–1 | Favours reactants | Reactants predominate | Weak acid/base equilibria |
| K << 10⁻³ | Far left (reactants) | Reaction barely occurs | Water autoionisation Kw |
Interactive Tools
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Related Terms
Chemical Equilibrium
Chemical equilibrium is the state in a reversible reaction where the rates of the forward and reverse reactions are equal, resulting in no net change in the concentrations of reactants and products over time. The system appears static but is actually dynamic — molecules continuously react in both directions at matching rates. The equilibrium state is quantified by an equilibrium constant (K), whose value depends only on temperature for a given reaction.
Kc (equilibrium)
Kc is the equilibrium constant expressed in terms of molar concentrations (mol/L) of reactants and products at equilibrium. Each concentration is raised to the power of its stoichiometric coefficient, and pure solids and pure liquids are excluded from the expression because their concentrations are constant. Kc is temperature-dependent and is the most commonly used form of the equilibrium constant in solution-phase and heterogeneous equilibria.
Kp (equilibrium)
Kp is the equilibrium constant expressed in terms of the partial pressures of gaseous reactants and products, each raised to the power of their stoichiometric coefficients. It is used exclusively for reactions involving gases and is related to Kc through the ideal gas equation, with the conversion factor depending on the change in moles of gas in the reaction. Kp is particularly useful in industrial gas-phase reactions such as the Haber process and the Contact process.
From Latin "aequus" (equal) and "constans" (standing firm). The concept was quantified by Guldberg and Waage in their 1864 Law of Mass Action; the notation K was widely adopted in the 20th century through IUPAC standardisation.