Estimate reaction spontaneity from thermodynamic inputs. Switch between common equations with consistent units. Download clean reports for labs and classes.
| Method | Inputs | ΔG (kJ/mol) | Interpretation |
|---|---|---|---|
| ΔG = ΔH − TΔS | ΔH = −120.5 kJ/mol, ΔS = −0.135 kJ/(mol·K), T = 298.15 K | −80.25 | Spontaneous at given temperature |
| ΔG° = −RT ln K | T = 298.15 K, K = 2500 | −19.34 | Favors products at equilibrium |
| ΔG = ΔG° + RT ln Q | ΔG° = −15.2 kJ/mol, T = 298.15 K, Q = 0.25 | −18.64 | Driving force increases when Q<1 |
Gibbs free energy change, ΔG, connects heat, disorder, and temperature into one decision metric for reaction direction. When ΔG is negative, the process is thermodynamically favorable under the stated conditions. When ΔG is positive, the forward direction is unfavorable unless conditions shift. In laboratory planning, ΔG helps prioritize which reactions merit optimization before investing in catalysts, separations, or scale‑up steps.
The relation ΔG = ΔH − TΔS shows why temperature can flip spontaneity. A reaction with positive ΔS gains a larger −TΔS term as temperature rises, often becoming more favorable at higher T. Conversely, when ΔS is negative, increasing T makes ΔG less favorable. Typical reporting uses ΔH in kJ/mol and ΔS in J/(mol·K), so consistent unit conversion is essential before multiplying by Kelvin.
For many systems, ΔG° is obtained from ΔG° = −RT ln K. Here, K is dimensionless and reflects equilibrium composition at a given temperature. Large K values yield strongly negative ΔG°, indicating product‑favored equilibrium. Small K values yield positive ΔG°, indicating reactant‑favored equilibrium. Because ln is sensitive when K is near 1, using sufficient numeric precision and temperature in Kelvin improves reliability.
Real mixtures rarely start at equilibrium. The adjustment ΔG = ΔG° + RT ln Q incorporates current activities via Q. If Q is less than K, ln(Q/K) is negative and the system tends to proceed forward. If Q exceeds K, the sign reverses and the reverse direction is favored. This perspective is useful for biochemical pathways, electrochemical cells, and gas‑phase syntheses where concentrations change during operation.
Professional reporting pairs ΔG with assumptions: temperature, standard state, and whether values are molar and dimensionless. Practical checks include confirming Kelvin conversion, verifying K and Q are positive, and reviewing the sign of ΔH and ΔS. Exporting results to CSV supports lab notebooks and audit trails, while a concise PDF snapshot helps communicate feasibility, sensitivity to temperature, and the expected driving force. For teaching, compare methods side by side to see how equilibrium, entropy, and composition each influence outcomes clearly.
A negative ΔG indicates the forward reaction is thermodynamically favorable under the specified temperature and composition. It does not guarantee fast rate; kinetics and activation energy still matter.
The equations use absolute temperature because entropy terms scale with thermal energy. Converting °C to K prevents incorrect magnitudes and sign changes when evaluating TΔS or RT ln terms.
In these relationships, K and Q are treated as dimensionless ratios of activities to standard states. Using concentrations or partial pressures is common when they approximate activities reasonably.
ΔG° applies to standard conditions, typically 1 bar for gases and 1 M for solutes. ΔG adjusts ΔG° to your current mixture using RT ln Q.
Accuracy depends on the quality of ΔH and ΔS data and whether they are valid over the temperature range. If values vary strongly with temperature, consider heat‑capacity corrections.
Report the method used, temperature, assumed standard state, and units. Include whether K or Q came from activities, concentrations, or pressures, and note significant figures consistent with inputs.
Important Note: All the Calculators listed in this site are for educational purpose only and we do not guarentee the accuracy of results. Please do consult with other sources as well.