Gibbs free energy tells you whether a reaction can happen on its own at a given temperature. If ΔG = ΔH − TΔS is negative the reaction is spontaneous; if it is positive it is not; if it is zero the system is at equilibrium.
Enter the enthalpy change in kJ/mol, the entropy change in J/(mol·K) and the temperature. Using standard values (ΔH°, ΔS°) also gives the equilibrium constant from ΔG° = −RT ln K. The defaults are for the Haber process, N₂ + 3H₂ → 2NH₃.
| Enthalpy change (ΔH) | -92.2 kJ/mol |
| Entropy change (ΔS) | -198.7 J/(mol·K) |
| Temperature (T) | 25 °C |
| ΔG (kJ/mol) | -32.9576 |
| Equilibrium constant K | 594,184 |
| Temperature (K) | 298.15 |
ΔG = -32.9576 kJ/mol, so making ammonia is spontaneous at room temperature.
Try this example →| Enthalpy change (ΔH) | -92.2 kJ/mol |
| Entropy change (ΔS) | -198.7 J/(mol·K) |
| Temperature (T) | 500 °C |
| ΔG (kJ/mol) | 61.4249 |
| Equilibrium constant K | 7.08214e-5 |
| Temperature (K) | 773.15 |
At 500 °C, ΔG = 61.4249 kJ/mol: the reaction is no longer spontaneous, which is why a balance of temperature and pressure is used.
Try this example →ΔG = ΔH − TΔS.
When ΔG is negative.
ΔH is usually in kJ/mol and ΔS in J/(mol·K), so ΔS must be converted to kJ to match.
ΔG° = −RT ln K. A negative ΔG° gives K > 1, favouring products.