Calculate equilibrium constant Kc, Kp, and reaction quotient Q. Determine equilibrium concentrations, predict reaction direction, and convert between Kc and Kp.
The equilibrium constant (K) is a dimensionless number that expresses the ratio of product concentrations to reactant concentrations at equilibrium, each raised to the power of their stoichiometric coefficients. It is one of the most powerful concepts in chemistry, allowing prediction of reaction direction, equilibrium compositions, and the effect of changing conditions.
For reactions in solution, we use Kc (concentrations in mol/L). For gas-phase reactions, Kp (partial pressures) is often more convenient. The two are related by Kp = Kc(RT)^Δn, where Δn is the change in moles of gas. The reaction quotient Q has the same form as K but uses current (non-equilibrium) concentrations — comparing Q to K reveals whether a reaction will shift forward or backward.
This calculator handles multiple scenarios: computing K from equilibrium concentrations, predicting equilibrium concentrations from K and initial values (ICE table), converting between Kc and Kp, and calculating Q to predict reaction direction. It also relates K to the standard Gibbs energy: ΔG° = -RT ln K.
Equilibrium calculations often involve solving polynomial equations (ICE tables) and careful unit management. This calculator automates these steps and provides the thermodynamic context (ΔG°) that connects equilibrium to energy. This equilibrium constant calculator helps you compare outcomes quickly and reduce avoidable mistakes when making day-to-day care decisions. Use the estimate as a planning baseline and confirm final decisions with a qualified professional when risk is high.
For aA + bB ⇌ cC + dD:\n\n Kc = [C]^c × [D]^d / ([A]^a × [B]^b)\n Kp = (P_C)^c × (P_D)^d / ((P_A)^a × (P_B)^b)\n Kp = Kc × (RT)^Δn where Δn = (c+d) - (a+b)\n ΔG° = -RT ln K\n Q has same form as K but at non-equilibrium conditions This keeps planning practical and lowers the chance of preventable errors.
Result: Kc = 0.211
For N₂O₄(g) ⇌ 2NO₂(g), Kc = [NO₂]²/[N₂O₄] = (0.0172)²/0.00140 = 2.958×10⁻⁴/1.40×10⁻³ = 0.211. This moderate K value means both reactants and products are present in significant amounts at equilibrium.
Chemical equilibrium is a dynamic state where the forward and reverse reactions occur at equal rates, resulting in constant concentrations over time. The equilibrium constant K quantifies this balance. It depends only on temperature, not on initial concentrations or the presence of a catalyst.
The ICE (Initial, Change, Equilibrium) table is the standard approach for computing equilibrium concentrations from K and initial conditions. Set up the table with initial concentrations, express changes in terms of x, and substitute into the K expression to solve for x. This often leads to quadratic (or higher) equations.
The Gibbs energy relationship ΔG° = -RT ln K is one of the most important equations in chemistry. It bridges the gap between thermodynamic tables (ΔG° values) and practical equilibrium predictions. At room temperature, ΔG° = -5.7 kJ/mol corresponds to K = 10, and ΔG° = -57 kJ/mol corresponds to K = 10¹⁰.
K >> 1 means the equilibrium strongly favors products. K << 1 means reactants are favored. K ≈ 1 means significant amounts of both are present.
Yes, K depends on temperature through the van't Hoff equation. For exothermic reactions, K decreases as temperature increases; for endothermic reactions, K increases.
When a system at equilibrium is disturbed, it shifts to partially counteract the disturbance. Adding reactant shifts toward products; increasing pressure shifts toward fewer moles of gas.
If Q < K, the reaction shifts right (toward products). If Q > K, it shifts left (toward reactants). If Q = K, the system is at equilibrium.
Use Kc for reactions in solution or when concentrations are given. Use Kp for gas-phase reactions when partial pressures are given. They're related by Kp = Kc(RT)^Δn.
No. Pure solids and pure liquids have activity = 1 and are omitted from the equilibrium expression. Only gases and dissolved species are included.