Equilibrium Constant (Keq) Calculator
Enter coefficients and equilibrium concentrations for up to two reactants and two products to compute Kc, and see which side of the reaction is favored.
Reviewed by the ToolNestr Editorial Team — July 2026
Leave a species' concentration blank to exclude it from the expression (useful for single-reactant or single-product reactions).
Reactants (denominator)
Products (numerator)
Two ideas that trip students up
1. A big Kc means a big product pile
The left pair shows a products-favored reaction (Kc >> 1): a small reactant pile, a large product pile. The right pair shows the opposite (Kc << 1): mostly reactants, very little product.
2. Kc spans many orders of magnitude
Three illustrative reactions side by side: a strongly reactant-favored one (Kc ≈ 0.01), a balanced one (Kc ≈ 1), and a strongly product-favored one (Kc ≈ 100) — bar height is shown on a compressed scale just to make all three visible together.
Equilibrium graphs
How it works
The core idea in one line: at equilibrium, the ratio of product concentrations to reactant concentrations — each raised to its stoichiometric coefficient — settles at a fixed number, Kc, for a given reaction at a given temperature.
Kc = [C]c[D]d / [A]a[B]b
products over reactants, each raised to its coefficient
Kc ≫ 1 → products favored
Kc ≪ 1 → reactants favored
Kc >> 1 means the numerator (products) dominates the ratio, so the reaction proceeds nearly to completion. Kc << 1 means the denominator (reactants) dominates, so very little product forms. A blank species field is treated as absent, so the expression naturally collapses for reactions with only one reactant or one product.
Worked example 1 — ammonia synthesis
Given: N₂ + 3 H₂ ⇌ 2 NH₃. At equilibrium: [N₂] = 0.50 M, [H₂] = 0.30 M, [NH₃] = 0.20 M.
Kc ≈ 2.96 is close to 1, so at this particular set of concentrations neither side is overwhelmingly favored.
Worked example 2 — simple 1:1 reaction
Given: A ⇌ B. At equilibrium: [A] = 0.10 M, [B] = 0.40 M. Both coefficients are 1, and there is no second reactant or product.
Kc = 4 > 1 means product B is somewhat favored at equilibrium.
Reading the size of Kc
The value of Kc tells you the position of equilibrium, not how fast it gets there.
| Kc value | Interpretation | Illustrative example |
|---|---|---|
| Kc ≫ 1 (e.g. > 10³) | Products strongly favored — reaction runs nearly to completion | Many combustion reactions |
| Kc ≈ 1 | Comparable amounts of reactants and products at equilibrium | Some gas-phase equilibria, e.g. certain isomerizations |
| Kc ≪ 1 (e.g. < 10⁻³) | Reactants strongly favored — very little product forms | Dissociation of weak acids in water |
These are conceptual bands, not fixed cutoffs — "large" and "small" are relative to the precision needed for the problem at hand.
Where equilibrium constants actually matter
🏭 Haber process (ammonia synthesis)
Industrial ammonia production (N₂ + 3H₂ ⇌ 2NH₃) is run at conditions chosen using Kc and Le Chatelier's principle to push the equilibrium toward more product while keeping the reaction rate practical.
🧬 Biochemical equilibria
Enzyme-catalyzed reactions, oxygen binding to hemoglobin, and metabolic pathways all involve equilibrium constants that determine how far a reaction proceeds inside cells under physiological concentrations.
🌫️ Environmental chemistry
Equilibrium constants describe how pollutants partition between air, water and soil, and how dissolved carbon dioxide equilibrates with carbonate species in oceans — central to modeling ocean acidification.
Common misconceptions
"Adding more reactant changes Keq."
Keq is constant at a fixed temperature. Adding reactant shifts the position of equilibrium — the system produces more product to partially counteract the change (Le Chatelier's principle) — but the same Kc expression is satisfied again with the new concentrations.
"A large Keq means the reaction is fast."
Keq describes the equilibrium position (how far the reaction goes), not the rate at which it gets there. A reaction can have a huge Keq and still be extremely slow without a catalyst.
"Keq is the same as Kp."
Kc uses concentrations and Kp uses partial pressures; they are numerically equal only when Δn = 0, and are otherwise related by Kp = Kc(RT)^Δn.
"If Kc = 1 the reaction hasn't started."
Kc = 1 simply means products and reactants happen to be present in comparable amounts at equilibrium — the reaction has fully reached equilibrium, it just is not lopsided toward either side.
Formula sources & further reading
The formulas here are standard, traceable to:
- • OpenStax, Chemistry 2e — Chapter 13, Fundamental Equilibrium Concepts (free, peer-reviewed). openstax.org
- • Brown, LeMay & Bursten, Chemistry: The Central Science — Chapter 15, Chemical Equilibrium.
- • Zumdahl & Zumdahl, Chemistry — Chapter 13, Chemical Equilibrium.
A blank reactant or product field is treated as absent (excluded from the expression). Results are rounded for display.
How to use this calculator
Enter reactants
Coefficient and equilibrium concentration for A, optionally B.
Enter products
Coefficient and equilibrium concentration for C, optionally D.
Read Kc
See the computed Kc and whether products or reactants are favored.
Related tools
Frequently asked questions
What is the equilibrium constant Keq?
Keq (often written Kc when concentrations are used) is the ratio of product concentrations to reactant concentrations, each raised to its stoichiometric coefficient, once a reversible reaction has reached equilibrium: Kc = [C]^c[D]^d / [A]^a[B]^b. It is a fixed number for a given reaction at a given temperature.
What does a large or small Keq mean?
Kc >> 1 means the numerator dominates — at equilibrium, products are strongly favored (the reaction proceeds nearly to completion). Kc << 1 means reactants are favored — very little product forms. Kc ≈ 1 means comparable amounts of reactants and products are present at equilibrium.
Does Keq change if I add more reactant or product?
No. Keq is constant at a given temperature. Adding more reactant or product shifts the position of equilibrium (Le Chatelier's principle) — the system reacts to partially offset the change — but once it re-settles, the same Kc expression is satisfied again with new concentrations.
What is the difference between Kc and Kp?
Kc uses molar concentrations (mol/L); Kp uses partial pressures, typically for gas-phase reactions. They are related by Kp = Kc(RT)^Δn, where Δn is the change in moles of gas (products − reactants) and R, T are the gas constant and absolute temperature. This calculator computes Kc from concentrations.
What happens if I only have one reactant or one product?
Leave the second reactant or second product field blank. This calculator treats a blank species as absent — its concentration and exponent are excluded from the expression (equivalent to a factor of 1), so the formula automatically collapses to match reactions with just one species on that side.