Equilibrium Constant Calculator
Determine the equilibrium constant for a chemical reaction by entering the concentrations of products and reactants with their stoichiometric coefficients. Supports Kc calculations for reactions with up to two products and two reactants.
Every reversible reaction reaches equilibrium — a state where the forward and reverse reactions proceed at equal rates and concentrations stop changing. The equilibrium constant K describes where the balance sits: large K means products dominate, small K means reactants dominate, K ≈ 1 means roughly equal amounts. Once you know K for a reaction (from a table or measurement), you can predict the concentrations at any starting point and answer the central question of practical chemistry: how much product can I actually expect?
This calculator computes K_c (concentration-based equilibrium constant) from the equilibrium concentrations of products and reactants. The general expression is K_c = [products]ⁿ / [reactants]ᵐ, with each concentration raised to its stoichiometric coefficient from the balanced equation. For a reaction aA + bB ⇌ cC + dD, K_c = [C]ᶜ[D]ᵈ / [A]ᵃ[B]ᵇ. Plug in concentrations and exponents and the calculator returns K.
Equilibrium constants run from very small (10⁻³⁰ for some unfavorable reactions) to very large (10²⁰ for strongly product-favored reactions), spanning more orders of magnitude than almost any other chemistry quantity. They're at the heart of pH calculations (K_w = 10⁻¹⁴ for water), solubility (K_sp), acid strength (K_a), buffer design, biochemistry (enzyme binding, protein folding), industrial chemistry, and environmental fate of pollutants. Mastering the K_c expression unlocks all of them.
Inputs
Leave 0 if only one product
Leave 0 if only one reactant
Results
Keq
5.000e+0
Favored
Products favored
ΔG°
-3.99 kJ/mol
Equilibrium Constant Results
| Parameter | Value |
|---|---|
| Product 1: [P1]^n1 | 0.5^1 = 5.0000e-1 |
| Product 2: [P2]^n2 | N/A |
| Reactant 1: [R1]^n1 | 0.1^1 = 1.0000e-1 |
| Reactant 2: [R2]^n2 | N/A |
| Keq | 5.0000e+0 |
| log(Keq) | 0.6990 |
| Favored Side | Products favored |
| ΔG° at 298 K | -3.9895 kJ/mol |
Formula
How to use this calculator
- Identify the balanced equation. Coefficients in the K expression come directly from this.
- Measure or look up equilibrium concentrations of each species. These are the values when the reaction has reached equilibrium (concentrations no longer changing).
- Enter products on the top, reactants on the bottom. The calculator returns K_c.
- For reactions involving only one product or only one reactant, set the other to 0 (the calculator ignores those terms).
- For solid or pure-liquid components, exclude them — set coefficient to 0 or leave concentration at default 1 M to make them inert in the formula.
- Compare to known K values to verify your measurements: K > 1 (products favored), K < 1 (reactants favored), K ≈ 1 (balanced).
Worked examples
Industrial ammonia synthesis (Haber process)
**Scenario:** N₂ + 3 H₂ ⇌ 2 NH₃. At 25 °C and equilibrium: [N₂] = 0.50 M, [H₂] = 1.5 M, [NH₃] = 0.30 M. Compute K_c. **Calculation:** K_c = [NH₃]² / ([N₂][H₂]³) = (0.30)² / (0.50 × 1.5³) = 0.090 / 1.6875 = 0.0533. **Result:** K_c = 0.053 at 25 °C — small, meaning reactants are favored. Industrial Haber-Bosch synthesis runs at 400–500 °C and 200 atm to shift equilibrium toward NH₃ (despite the fact that the high temperature actually reduces K — pressure dominates).
Weak acid dissociation
**Scenario:** Acetic acid in water: CH₃COOH ⇌ H⁺ + CH₃COO⁻. At equilibrium in a 0.1 M solution: [H⁺] = 1.3 × 10⁻³ M, [CH₃COO⁻] = 1.3 × 10⁻³ M, [CH₃COOH] ≈ 0.1 M (since only ~1% dissociates). **Calculation:** K_a = [H⁺][CH₃COO⁻] / [CH₃COOH] = (1.3 × 10⁻³)² / 0.1 = 1.69 × 10⁻⁶ / 0.1 = 1.69 × 10⁻⁵. **Result:** K_a ≈ 1.7 × 10⁻⁵, pK_a = 4.76 — matches the literature value for acetic acid. This explains why vinegar (5% acetic acid) has pH ~2.5, not pH 0 (which would require complete dissociation like HCl).
Predicting reaction direction with Q vs K
**Scenario:** Reaction A ⇌ B has K_c = 5.0 at 25 °C. You mix [A] = 0.20 M and [B] = 0.40 M. Will the reaction proceed forward or backward? **Calculation:** Reaction quotient Q = [B] / [A] = 0.40 / 0.20 = 2.0. Compare to K = 5.0: Q < K, so reaction proceeds forward (toward more product). Equilibrium will continue forming B from A until [B]/[A] = 5.0. **Result:** Q < K → reaction proceeds forward. Final equilibrium: if total = 0.60 M, then [A]_eq + [B]_eq = 0.60 and [B]_eq/[A]_eq = 5.0 → [A]_eq = 0.10 M, [B]_eq = 0.50 M. So 0.10 M of A will convert to B as the system equilibrates.
When to use this calculator
**Use equilibrium constant math for:**
- **Predicting reaction extent**: K large → goes nearly to completion; K small → barely any product forms. - **Designing industrial processes**: Haber, contact (H₂SO₄), Ostwald (HNO₃), all use equilibrium analysis to choose T, P, and recycle strategies. - **Buffer design**: K_a determines the useful pH range of a buffer; Henderson-Hasselbalch is a rearrangement of K_a expression. - **Solubility predictions**: K_sp tells you the concentration at which precipitate forms. - **Acid/base strength**: pK_a is just −log₁₀(K_a); strong vs weak distinction comes from K. - **Biochemistry**: enzyme binding (K_d), protein folding (K_unfold), receptor-ligand interactions. - **Environmental chemistry**: partitioning of pollutants between water/air/soil phases, all equilibrium-driven. - **Geochemistry**: mineral solubility, carbonate equilibria controlling ocean pH.
**Practical guidelines for using K:**
- **K > 1000**: reaction is essentially irreversible (goes to completion). - **K = 10 to 1000**: products strongly favored; high but not complete yield. - **K = 0.1 to 10**: mixture of products and reactants at equilibrium. - **K = 10⁻³ to 0.1**: reactants strongly favored; little product forms. - **K < 10⁻³**: essentially no reaction occurs.
**Le Chatelier — practical levers:**
- **Increase reactant concentration**: shifts toward more product. - **Remove product as formed**: shifts further toward product (Haber recycles unreacted N₂/H₂). - **Increase pressure** (gas reactions with fewer moles of gas product): shifts toward product. - **Increase temperature for endothermic reactions**: increases K. - **Decrease temperature for exothermic reactions**: increases K (but slower reaction). - **Catalyst**: speeds equilibration but doesn't change K.
**K depends on temperature:**
K_eq(T) is a strong function of temperature via van't Hoff equation:
ln(K₂/K₁) = −ΔH°/R × (1/T₂ − 1/T₁)
So K reported at one temperature is different at another. Always check temperature when comparing K values.
Common mistakes to avoid
- Including pure solids and pure liquids in the K expression. They have constant "concentration" (activity = 1) and are omitted.
- Using initial concentrations instead of equilibrium concentrations. K is only defined when the system has reached equilibrium.
- Skipping stoichiometric coefficients. [A]² is very different from [A] when [A] = 0.5 (0.25 vs 0.5).
- Confusing K with rate constants. K (equilibrium) describes the ratio at equilibrium; k (rate constant) describes how fast the reaction proceeds. They're different.
- Using K_c when the problem requires K_p (or vice versa) for gas-phase reactions. They're related by K_p = K_c × (RT)^Δn.
- Forgetting that K depends on temperature. A K value from a 25 °C table is wrong at 100 °C.
- Confusing K with the reaction quotient Q. Q uses any-time concentrations; K uses specifically equilibrium concentrations. Q tells you which direction the system is shifting toward equilibrium.