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Equilibrium Constant Calculator

Chemistry

Calculate the equilibrium constant Kc from equilibrium concentrations of reactants and products for a reversible reaction. Also compute ΔG° from Kc at 25°C.

0.5 mol/L
mol/L
2
0.2 mol/L
mol/L
1

Equilibrium Constant (Kc)

1.25
log Kc
0.097
ΔG° at 25°C (kJ/mol)
-0.553

This calculator computes your Equilibrium Constant (Kc), log Kc, ΔG° at 25°C (kJ/mol) from the values you enter.

Inputs
Product Concentration [P] (mol/L)Stoichiometric Coefficient of ProductReactant Concentration [R] (mol/L)Stoichiometric Coefficient of Reactant
Outputs
Equilibrium Constant (Kc)log KcΔG° at 25°C (kJ/mol)

What is a Equilibrium Constant?

The Equilibrium Constant Calculator computes Kc — the dimensionless equilibrium constant — from the equilibrium concentrations and stoichiometric coefficients of one product and one reactant species in a reversible chemical reaction. It also calculates log Kc (for use in thermodynamic relationships) and the standard Gibbs free energy change ΔG° at 25°C, connecting chemical equilibrium directly to thermodynamics.

Kc is the central quantity in chemical equilibrium analysis. For any reversible reaction at a fixed temperature, the ratio of product concentrations to reactant concentrations (each raised to their stoichiometric coefficients) reaches a constant value at equilibrium — this is Kc. A Kc much greater than 1 means products predominate at equilibrium; a Kc much less than 1 means reactants predominate. Temperature is the only variable that changes Kc — adding more reactant, removing product, or changing pressure shifts the position of equilibrium but not the value of Kc itself.

The relationship between Kc and thermodynamics is direct: ΔG° = −RT ln(Kc). A reaction with Kc >> 1 has a large negative ΔG°, meaning products are thermodynamically much more stable than reactants. This tool displays ΔG° at 25°C (T = 298.15 K) using R = 8.314 J/(mol·K). For ΔG° calculations at other temperatures or using enthalpy and entropy directly, use the Gibbs Free Energy Calculator.

For reactions not yet at equilibrium, the reaction quotient Q has the same form as Kc but uses current (non-equilibrium) concentrations. Comparing Q to Kc predicts the direction of spontaneous reaction.

How to use this Equilibrium Constant calculator

  1. Write the balanced equation for your reversible reaction and identify the equilibrium concentrations (in mol/L) of all species from your data or ICE table.
  2. Enter the equilibrium concentration of the product species in the Product Concentration [P] (mol/L) field. For multi-product reactions, calculate the numerator manually: [C]^c × [D]^d, and enter the result as a single equivalent concentration (with coefficient 1) if using this calculator for the full product term.
  3. Enter the stoichiometric coefficient of the product in the Stoichiometric Coefficient of Product field.
  4. Enter the equilibrium concentration of the reactant in the Reactant Concentration [R] (mol/L) field. Similarly, if there are multiple reactants, combine them manually.
  5. Enter the stoichiometric coefficient of the reactant in Stoichiometric Coefficient of Reactant.
  6. Read Kc — note whether it is greater or less than 1, and by how many orders of magnitude. Read ΔG° (kJ/mol) to confirm the thermodynamic spontaneity direction.

Formula & Methodology

Kc expression (single product, single reactant):

Kc = [P]^nP / [R]^nR

General Kc expression:

Kc = [C]^c × [D]^d / ([A]^a × [B]^b)      for aA + bB ⇌ cC + dD

Derived outputs:

log Kc = log₁₀(Kc) ΔG° (kJ/mol) = −R × T × ln(Kc) / 1000               where R = 8.314 J/(mol·K), T = 298.15 K

Worked example — hydrogen iodide equilibrium:

Reaction: H₂(g) + I₂(g) ⇌ 2 HI(g) at 445°C

Equilibrium concentrations measured: [H₂] = 0.107 mol/L, [I₂] = 0.107 mol/L, [HI] = 0.786 mol/L.

Kc = [HI]² / ([H₂]¹ × [I₂]¹)    = (0.786)² / (0.107 × 0.107)    = 0.618 / 0.01145    = 53.97  log Kc = log(53.97) = 1.732  ΔG° at 25°C = −(8.314)(298.15) ln(53.97) / 1000              = −2478.8 × 3.988 / 1000              = −9.88 kJ/mol

Kc = 54 is greater than 1, confirming that HI is favoured at equilibrium at 445°C. The negative ΔG° at 25°C indicates that HI is also thermodynamically favoured at room temperature, though the specific Kc value at 25°C would differ from the one measured at 445°C.

Frequently Asked Questions

The equilibrium constant Kc is a dimensionless number that quantifies the ratio of product concentrations to reactant concentrations at equilibrium for a reversible reaction at a given temperature. Each concentration is raised to the power of its stoichiometric coefficient from the balanced equation. A large Kc (much greater than 1) means the equilibrium lies toward the products; a small Kc (much less than 1) means it lies toward the reactants. Kc is temperature-dependent — it changes when temperature changes but is unaffected by concentration changes, pressure changes, or catalysts.
For a reaction aA + bB ⇌ cC + dD, the equilibrium constant expression is Kc = [C]^c × [D]^d / ([A]^a × [B]^b), where square brackets denote molar concentrations (mol/L) at equilibrium and the exponents are the stoichiometric coefficients from the balanced equation. Pure solids and pure liquids (including water in dilute aqueous solutions) are omitted from the Kc expression — their concentrations are constant and absorbed into the Kc value.
Kc uses molar concentrations (mol/L) in its expression; Kp uses partial pressures (usually in atm or Pa) of gaseous reactants and products. For reactions involving gases, both are defined and are related by Kp = Kc × (RT)^Δn, where R is the gas constant (0.0821 L·atm/mol·K), T is the absolute temperature in Kelvin, and Δn is the change in moles of gas (moles of gaseous products minus moles of gaseous reactants). For reactions with no change in moles of gas (Δn = 0), Kc and Kp are numerically equal.
The standard Gibbs free energy change ΔG° is related to Kc by ΔG° = −RT ln(Kc). A negative ΔG° means Kc > 1 — the reaction is thermodynamically favoured in the forward direction and products are favoured at equilibrium. A positive ΔG° means Kc < 1 — the reaction is thermodynamically unfavoured in the forward direction and reactants predominate at equilibrium. A ΔG° of zero means Kc = 1 — neither reactants nor products are favoured. This relationship connects thermodynamics and equilibrium, linking the [Gibbs Free Energy Calculator](/gibbs-free-energy-calculator/) to Kc.
The reaction quotient Q has the same mathematical form as Kc — products over reactants with concentration exponents — but uses current concentrations instead of equilibrium concentrations. Comparing Q to Kc predicts which direction the reaction will proceed: if Q < Kc, the reaction proceeds forward (toward products); if Q > Kc, it proceeds in reverse (toward reactants); if Q = Kc, the system is already at equilibrium. The Reaction Quotient Calculator helps compute Q and compare it to Kc.
Kc is temperature-dependent — it is a constant only at a fixed temperature. For an exothermic reaction (negative ΔH°), increasing temperature shifts the equilibrium toward reactants and decreases Kc (Le Chatelier's principle). For an endothermic reaction (positive ΔH°), increasing temperature shifts equilibrium toward products and increases Kc. The quantitative relationship is given by the van't Hoff equation: ln(K₂/K₁) = −ΔH°/R × (1/T₂ − 1/T₁), which relates Kc values at two temperatures T₁ and T₂.
Enter the equilibrium concentration of products in the 'Product Concentration' field (in mol/L) and the product's stoichiometric coefficient in its coefficient field. Enter the equilibrium concentration of reactants in the 'Reactant Concentration' field and its coefficient. The calculator returns Kc using the formula Kc = [P]^nP / [R]^nR, along with log Kc and the standard Gibbs free energy change ΔG° at 25°C.
The current calculator handles one product species and one reactant species — it covers the most common single-step equilibrium problems taught in NCERT, JEE, and NEET chemistry. For reactions with multiple products or reactants, apply the formula manually: Kc = ([P1]^p1 × [P2]^p2) / ([R1]^r1 × [R2]^r2). Multiply together each species' concentration raised to its coefficient on each side, then divide products by reactants.
Yes — equilibrium constants govern the design of major industrial chemical processes. The Haber process for ammonia synthesis (N₂ + 3H₂ ⇌ 2NH₃), critical for fertiliser production, operates at conditions selected to balance Kc (temperature, pressure, catalyst) for maximum yield. India is a major fertiliser producer, and plants at IFFCO, NFL, and Rashtriya Chemicals operate these equilibrium-governed processes at scale. Understanding Kc enables process engineers to optimise temperature and pressure conditions to shift equilibrium toward the desired product.
Kc values can span many orders of magnitude — from 10⁻³⁰ for reactions strongly favouring reactants to 10⁴⁰ for reactions that essentially go to completion. Log Kc compresses this range into a linear scale: log Kc = 0 means Kc = 1 (balanced equilibrium); log Kc = 10 means Kc = 10¹⁰ (strongly product-favoured). Log Kc is also the form that appears directly in the ΔG° equation (ΔG° = −2.303 RT × log Kc) and in electrochemistry (related to standard cell potential via the Nernst equation).
For the reaction 2 H₂(g) + O₂(g) ⇌ 2 H₂O(g) at 25°C, Kc is approximately 3 × 10⁸⁰ — an astronomically large value indicating that the equilibrium lies overwhelmingly toward water at room temperature. This is reflected in the very negative ΔG° (≈ −457 kJ/mol). In practice, this reaction does not proceed measurably at 25°C without a spark or catalyst because the activation energy is very high — Kc tells you only about thermodynamic favourability, not reaction rate.