Calculate concentration equilibrium constants (Kc), gas partial pressure constants (Kp), compare reaction quotients (Q) against K to predict directional shifts, and solve standard Gibbs free energy (ΔG°) with our free Equilibrium Constant Calculator.
Scientific: 5.476 × 10^1 • Forward Products Strongly Favored
Spontaneous
K > 1
Δn = 0
25.0 °C
[C]^c · [D]^d
[A]^a · [B]^b
In chemical thermodynamics, the Equilibrium Constant (\(K\)) quantifies the dynamic ratio between products and reactants when forward and reverse reaction rates become exactly equal:
The Law of Mass Action dictates that for any reversible process, the equilibrium constant depends solely on temperature and is invariant with changing initial concentrations. Comparing the instantaneous reaction quotient \(Q\) against \(K\) determines the spontaneous shift direction predicted by Le Chatelier's Principle.
Distinguishes initial non-equilibrium quotient values from true dynamic equilibrium constants to forecast Le Chatelier reaction shifts.
Computes polynomial exponents accurately without compounding manual multiplication errors.
Derives standard thermodynamic free energy changes \(\Delta G^\circ = -RT \ln K\) in \(\text{kJ/mol}\) with spontaneity indications.
Calculate equilibrium constants using solution molarity concentrations (\(\text{mol/L}\)) or gas-phase partial pressures (\(\text{atm/bar}\)) seamlessly.
Compares \(K\) against threshold values to diagnose whether products or reactants predominate at equilibrium.
Computes \(\Delta G^\circ = -RT \ln K\) at any specified temperature in Kelvin or Celsius to verify chemical spontaneity.
Generates beautifully typeset mass action fractions, exponent breakdowns, and thermodynamic substitutions in real time.
| Condition | Direction of Shift | Thermodynamic Free Energy (\(\Delta G\)) | System Behavior |
|---|---|---|---|
| \(Q < K\) | Forward Shift (\(\longrightarrow\)) | \(\Delta G < 0\) (Spontaneous Forward) | Reactants convert to additional products to reach equilibrium |
| \(Q = K\) | Dynamic Equilibrium (\(\rightleftharpoons\)) | \(\Delta G = 0\) (No Driving Force) | Forward rate equals reverse rate; macroscopic concentrations constant |
| \(Q > K\) | Reverse Shift (\(\longleftarrow\)) | \(\Delta G > 0\) (Spontaneous Reverse) | Excess products decompose back into reactants |
Choose an equilibrium preset or select between molar concentrations (\(K_c\)) and partial pressures (\(K_p\)).
Input measured equilibrium reactant and product values along with reaction temperature in Kelvin or Celsius.
Click "Calculate Equilibrium Constant" to evaluate \(K\), standard Gibbs energy \(\Delta G^\circ\), and full KaTeX derivations.
Review product favorability (\(K > 1\)), copy KaTeX steps, and integrate results into physical chemistry lab analyses.
Problem: In a sealed \(1.00\text{ L}\) reactor at \(448\text{ }^\circ\text{C}\) (\(721.15\text{ K}\)), the equilibrium concentrations for \(\text{H}_2(g) + \text{I}_2(g) \rightleftharpoons 2\text{HI}(g)\) are \([\text{H}_2] = 0.20\text{ M}\), \([\text{I}_2] = 0.20\text{ M}\), and \([\text{HI}] = 1.48\text{ M}\). Calculate \(K_c\) and determine standard Gibbs free energy \(\Delta G^\circ\).
1. Law of Mass Action:
2. Standard Gibbs Free Energy \(\Delta G^\circ\):
Problem: Acetic acid and ethanol react to produce ethyl acetate and water: \(\text{CH}_3\text{COOH} + \text{C}_2\text{H}_5\text{OH} \rightleftharpoons \text{CH}_3\text{COOC}_2\text{H}_5 + \text{H}_2\text{O}\). Equilibrium concentrations are \([\text{Acid}] = 0.33\text{ M}\), \([\text{Alcohol}] = 0.33\text{ M}\), \([\text{Ester}] = 0.67\text{ M}\), and \([\text{Water}] = 0.67\text{ M}\).
1. Equilibrium Constant: \(K_c = \frac{0.67 \times 0.67}{0.33 \times 0.33} = \frac{0.4489}{0.1089} = 4.12\)
2. Thermodynamic Free Energy: \(\Delta G^\circ = -(8.314) \times 298.15 \times \ln(4.12) = -3.51\text{ kJ/mol}\)
Substituting initial starting concentrations into \(K\) yields the reaction quotient (\(Q\)), not the equilibrium constant. Only values measured after concentration plateauing are valid in \(K\).
In dilute aqueous acid-base reactions, water functions as the solvent with an invariant activity of \(1.0\). It must be excluded from the denominator.
Catalysts speed up both forward and reverse reaction rates equally by lowering activation energy (\(E_a\)). They accelerate the time to reach equilibrium without changing the numerical value of \(K\).
\(K\) is a thermodynamic constant at a fixed temperature. Modifying reactor temperature shifts \(K\) according to the sign of the reaction enthalpy (\(\Delta H^\circ\)) via the van 't Hoff equation.
Human physiological blood pH is tightly regulated around \(7.40\) by the carbonic acid equilibrium: \(\text{CO}_2(aq) + \text{H}_2\text{O}(l) \rightleftharpoons \text{H}_2\text{CO}_3(aq) \rightleftharpoons \text{H}^+(aq) + \text{HCO}_3^-(aq)\) with combined \(K_a \approx 7.9 \times 10^{-7}\). Rapid respiration expels \(\text{CO}_2\) to compensate for metabolic acidosis.
Atmospheric carbon dioxide absorption shifts marine carbonate equilibria: \(\text{CO}_2 + \text{H}_2\text{O} + \text{CO}_3^{2-} \rightleftharpoons 2\text{HCO}_3^-\). Lowering ocean \([\text{CO}_3^{2-}]\) reduces the saturation state (\(\Omega\)) of aragonite, impeding coral reef calcification.
Authoritative answers to common questions regarding chemical equilibrium constants and Le Chatelier shifts.
