Equilibrium Constant Calculator
Calculates the equilibrium constant Kc from product and reactant concentrations for a simplified 1:1 stoichiometry reaction. Use it in general chemistry to predict reaction favorability.
About this calculator
The equilibrium constant Kc quantifies the ratio of product concentrations to reactant concentrations at equilibrium, with each term raised to the power of its stoichiometric coefficient. For a general reaction aA ⇌ bB, Kc = [B]ᵇ / [A]ᵃ. This calculator implements a simplified form for 1:2 stoichiometry: Kc = [products]² / [reactants]². A large Kc (≫ 1) means equilibrium lies toward products; a small Kc (≪ 1) means reactants predominate. Kp relates to Kc through Kp = Kc × (RT)^Δn, where R = 0.08206 L·atm/mol·K, T is temperature in Kelvin, and Δn is the change in moles of gas between products and reactants. Kc and Kp are equal when Δn = 0.
How to use
Suppose at equilibrium you measure [NO₂] = 0.200 M (product) and [N₂O₄] = 0.100 M (reactant) for the reaction N₂O₄ ⇌ 2 NO₂. Enter Product Concentration = 0.200 M, Reactant Concentration = 0.100 M. Kc = (0.200)² / (0.100)² = 0.0400 / 0.0100 = 4.00. Since Kc > 1, equilibrium favors NO₂ formation. If temperature is 298 K and Δn = +1, then Kp = 4.00 × (0.08206 × 298)¹ = 4.00 × 24.45 = 97.8 atm.
Frequently asked questions
What does a large or small equilibrium constant tell you about a reaction?
A Kc much greater than 1 (e.g., 10⁵) indicates the reaction strongly favors products at equilibrium — essentially going to completion. A Kc much less than 1 (e.g., 10⁻⁵) means reactants predominate and very little product forms. A Kc near 1 signals that both reactants and products are present in comparable amounts at equilibrium. These magnitudes guide chemists in deciding whether a reaction is synthetically useful or requires special conditions such as excess reagent or product removal to drive conversion.
How is Kp different from Kc and when should I use each one?
Kc uses molar concentrations (mol/L) while Kp uses partial pressures (atm or Pa), making Kp more convenient for gas-phase reactions measured with pressure gauges. They are related by Kp = Kc(RT)^Δn, where Δn is the change in moles of gaseous species from reactants to products. When Δn = 0 (equal moles of gas on both sides), Kp = Kc. For reactions involving solids or liquids alongside gases, only the gas-phase species appear in the Kp expression, because their 'concentrations' are incorporated into the constant.
Why does the equilibrium constant change with temperature but not with concentration or pressure?
The equilibrium constant is a thermodynamic quantity tied to the Gibbs free energy change of the reaction: ΔG° = −RT ln K. Because ΔG° depends on temperature through the enthalpy and entropy terms (ΔG° = ΔH° − TΔS°), K changes when temperature changes. Adding more reactant or product, or changing pressure, shifts the position of equilibrium (described by Le Chatelier's principle) but does not change the value of K itself — the system simply adjusts its concentrations until the same ratio is restored. Only temperature alters K.