Equilibrium Constant Calculator
Compute the equilibrium constant Kc from measured product and reactant concentrations for a reversible reaction. Use it in general chemistry or physical chemistry to analyze reaction equilibria.
Last updated: May 2026
About this calculator
For a generic reaction aA + bB ⇌ cC + dD, the equilibrium constant Kc is defined as: Kc = [C]^c × [D]^d / ([A]^a × [B]^b), where square brackets denote molar concentrations at equilibrium and the exponents match the stoichiometric coefficients. This calculator uses two products and two reactants each raised to the first power: Kc = ([P1] × [P2]) / ([R1] × [R2]). A large Kc (≫1) means products are favored; a small Kc (≪1) means reactants predominate. Kc changes with temperature but is unaffected by pressure or added catalysts. To link Kc to thermodynamics, use ΔG° = −RT ln Kc separately with the actual standard Gibbs energy change for your reaction.
How to use
Consider a reaction where at equilibrium [P1] = 0.5 M, [P2] = 0.3 M, [R1] = 0.1 M, [R2] = 0.2 M. Kc = (0.5 × 0.3) / (0.1 × 0.2) = 0.15 / 0.02 = 7.5. Enter those four concentration values. Since Kc > 1, products are favored at equilibrium. (Kc is defined directly from the equilibrium concentrations; to relate it to Gibbs energy, use ΔG° = −RT ln Kc separately.)
Frequently asked questions
What does a large equilibrium constant Kc tell you about a reaction?
A large Kc value (much greater than 1) indicates that at equilibrium the reaction mixture contains far more products than reactants—the reaction strongly favors the forward direction. For example, Kc = 10⁶ means the reaction is essentially complete under normal conditions. Conversely, Kc ≪ 1 means the reverse reaction is favored. A Kc near 1 indicates a mixture of significant amounts of both reactants and products at equilibrium.
How does temperature affect the equilibrium constant Kc?
Kc is temperature-dependent because the Gibbs energy change (ΔG° = −RT ln Kc) varies with T. For an exothermic reaction, increasing temperature shifts the equilibrium toward reactants, lowering Kc. For an endothermic reaction, raising temperature shifts equilibrium toward products, increasing Kc. This relationship is quantified by the van't Hoff equation: d(ln K)/dT = ΔH°/RT². Unlike concentration or pressure changes, a temperature change is the only factor that actually changes the value of Kc.
What is the difference between Kc and Kp for gas-phase equilibria?
Kc is expressed in terms of molar concentrations (mol/L), while Kp uses partial pressures (atm or Pa) of gaseous species. They are related by Kp = Kc × (RT)^Δn, where Δn is the change in moles of gas (moles of gaseous products minus moles of gaseous reactants) and R = 0.08206 L·atm/mol·K. When Δn = 0, Kp equals Kc. For reactions involving only solids or liquids, neither Kc nor Kp includes those pure-phase species because their activities are defined as 1.