Equilibrium Constant Calculator

At a glance:The equilibrium constant (concentration equilibrium constant, symbol Kc) is a quantitative measure of the relation between product and reactant concentrations when a reversible chemical reaction reaches equilibrium. Many reactions are reversible — the forward and reverse reactions proceed simultaneously; when their rates become equal, the species concentrations no longer change with time and the system reaches 'chemical equilibrium'. At that point, by the 'law of mass action', for a general reversible reaction aA + bB ⇌ cC + dD the concentration equilibrium constant is Kc = ([C]^c · [D]^d) / ([A]^a · [B]^b), i.e. 'the product of the products' equilibrium concentrations each raised to its stoichiometric power' divided by 'that of the reactants'. The square brackets [ ] denote the equilibrium molar concentration (mol/L), and the superscripts a, b, c, d are the stoichiometric coefficients of the respective species in the balanced equation. Using this tool's default (hydrogen–iodine reaction): H₂ + I₂ ⇌ 2HI, at equilibrium [HI] = 0.8 mol/L, [H₂] = [I₂] = 0.1 mol/L, so Kc = [HI]² / ([H₂][I₂]) = 0.8² / (0.1 × 0.1) = 0.64 / 0.01 = 64. The magnitude of Kc tells us 'which side the equilibrium lies on' and how completely the reaction proceeds: when Kc ≫ 1 (very large) the equilibrium strongly favours products and the reaction is nearly complete, with the forward reaction as the main direction; when Kc ≪ 1 (very small) the equilibrium favours reactants and the reaction barely proceeds; when Kc ≈ 1 the forward and reverse reactions are evenly matched and both reactants and products are present in appreciable amounts. Importantly, the equilibrium constant depends only on 'temperature' — for the same reaction, as long as the temperature is unchanged, no matter the initial concentrations or how equilibrium is reached, Kc stays the same fixed value. Changing concentration or pressure shifts the system to a new equilibrium (Le Chatelier's principle) but leaves Kc unchanged as long as temperature is constant; only changing temperature changes Kc (for an exothermic reaction, raising temperature lowers Kc; for an endothermic reaction, raising temperature raises Kc). Key rules for writing the equilibrium expression: first, pure solids and pure liquids (including water as the solvent) are omitted because their 'activity' is taken as 1; only gases and solutes in solution are included. Second, the coefficients become powers of the concentrations. Third, Kc uses concentrations; if gas partial pressures are used it becomes Kp, and the two are related by Kp = Kc(RT)^Δn. Applications of the equilibrium constant: judging how complete a reaction is, predicting the direction of shift (by comparing the reaction quotient Q with K), computing equilibrium concentrations, and linking to thermodynamics (ΔG° = −RT ln K). Using this calculator: it is designed as the general form aA + bB ⇌ cC + dD, with at most two reactants and two products. If a species is absent (e.g. only one product), simply 'set its coefficient to 0' — a term with coefficient 0 is counted as 1 and does not affect the product. The concentrations entered must be 'equilibrium' concentrations, not initial ones. The product of all reactant terms must not be 0 (otherwise the denominator is zero and Kc is undefined; this tool returns 0). In short, this calculator lets you quickly obtain the concentration equilibrium constant Kc from equilibrium concentrations and stoichiometric coefficients — the core tool for learning chemical equilibrium.