Equilibrium Constant Calculator

About the Equilibrium Constant Calculator

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.

When This Page Helps

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.

How to Use the Inputs

1. Select the calculation mode: compute K, ICE table, Kc↔Kp conversion, or Q vs K.
2. Enter equilibrium concentrations or pressures of reactants and products.
3. Enter stoichiometric coefficients for each species.
4. For ICE table mode, enter initial concentrations and K.
5. Select a preset for common equilibrium reactions.
6. Review K, Q, ΔG°, and the equilibrium position assessment.
7. Use the reference table to compare K values for different reactions.

Example Calculation

Tips & Best Practices

• K is dimensionless by convention when using activities, but Kc and Kp have implied units based on concentrations/pressures.
• Multiplying a reaction by a factor n raises K to the nth power: K_new = K^n.
• Reversing a reaction inverts K: K_reverse = 1/K_forward.
• For very large K (>10⁶), the reaction essentially goes to completion.
• The temperature dependence of K is given by the van't Hoff equation: ln(K₂/K₁) = -ΔH°/R × (1/T₂ - 1/T₁).
• ΔG° = -RT ln K connects thermodynamics to equilibrium: negative ΔG° means K > 1.

Understanding Chemical 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 Table Method

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.

Relationship Between K and Thermodynamics

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.

Frequently Asked Questions

• What does a large K mean?

  K >> 1 means the equilibrium strongly favors products. K << 1 means reactants are favored. K ≈ 1 means significant amounts of both are present.

• Does K change with temperature?

  Yes, K depends on temperature through the van't Hoff equation. For exothermic reactions, K decreases as temperature increases; for endothermic reactions, K increases.

• What is Le Chatelier's principle?

  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.

• How do Q and K compare?

  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.

• When do I use Kc vs Kp?

  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.

• Are solids and liquids included in K?

  No. Pure solids and pure liquids have activity = 1 and are omitted from the equilibrium expression. Only gases and dissolved species are included.