Calculate Gibbs free energy (ΔG) from enthalpy and entropy, determine spontaneity, and find equilibrium temperatures. Supports ΔG°rxn from formation data and ΔG = ΔH - TΔS.
Gibbs free energy (G) is the thermodynamic potential that determines whether a process will occur spontaneously at constant temperature and pressure. The change in Gibbs energy, ΔG = ΔH - TΔS, combines the enthalpy change (ΔH) and entropy change (ΔS) into a single criterion: if ΔG < 0, the process is spontaneous; if ΔG > 0, it is non-spontaneous; and if ΔG = 0, the system is at equilibrium.
Named after Josiah Willard Gibbs, this quantity is arguably the most important in chemical thermodynamics. It determines reaction feasibility, phase equilibria, electrochemical cell potentials (ΔG = -nFE°), and the equilibrium constant (ΔG° = -RT ln K). Understanding Gibbs energy allows you to predict which reactions will occur, at what temperatures they become favorable, and how much useful work they can perform.
This calculator computes ΔG from ΔH and ΔS at any temperature, determines the crossover temperature where ΔG changes sign, and relates ΔG° to the equilibrium constant K. It also calculates ΔG°rxn from standard free energies of formation.
Gibbs energy calculations require careful attention to units (kJ vs J, K vs °C) and signs. This calculator handles all conversions, determines spontaneity, computes the crossover temperature, and provides the corresponding K value. This gibbs free energy calculator helps you compare outcomes quickly and reduce avoidable mistakes when making day-to-day care decisions. Use the estimate as a planning baseline and confirm final decisions with a qualified professional when risk is high.
ΔG = ΔH - TΔS\n\nSpontaneity criterion:\n ΔG < 0 → spontaneous\n ΔG > 0 → non-spontaneous\n ΔG = 0 → equilibrium\n\nCrossover temperature: T_eq = ΔH / ΔS\n\nRelation to K: ΔG° = -RT ln K → K = exp(-ΔG°/RT)\n\nΔG°rxn = Σ n·ΔGf°(products) - Σ n·ΔGf°(reactants) This keeps planning practical and lowers the chance of preventable errors.
Result: ΔG = -32.98 kJ/mol (spontaneous)
For N₂ + 3H₂ → 2NH₃: ΔG = -92.2 - (298)(-0.1987) = -92.2 + 59.2 = -33.0 kJ/mol. The reaction is spontaneous at 298 K. The crossover temperature is 92200/198.7 = 464 K — above this, entropy dominates and the reaction becomes non-spontaneous.
The sign of ΔG = ΔH - TΔS depends on the signs of ΔH and ΔS. Case 1: ΔH < 0, ΔS > 0 — always spontaneous (exothermic + more disorder). Case 2: ΔH > 0, ΔS < 0 — never spontaneous. Case 3: ΔH < 0, ΔS < 0 — spontaneous at low T (enthalpy-driven). Case 4: ΔH > 0, ΔS > 0 — spontaneous at high T (entropy-driven, like CaCO₃ decomposition).
In biology, unfavorable reactions (ΔG > 0) are driven by coupling them with favorable ones. ATP hydrolysis (ΔG = -30.5 kJ/mol) drives countless otherwise non-spontaneous biochemical processes. The overall ΔG for coupled reactions is the sum of individual ΔG values.
The crossover temperature T_eq = ΔH/ΔS determines phase transition temperatures (melting, boiling) and reaction feasibility boundaries. Phase diagrams are essentially maps of where ΔG changes sign for different phases as functions of T and P.
Gibbs free energy (G) is a thermodynamic state function that predicts whether a process will occur spontaneously at constant T and P. ΔG < 0 means the process is spontaneous.
ΔG = 0 means the system is at equilibrium. The forward and reverse processes occur at equal rates, and there is no net change.
Temperature determines the relative importance of ΔH and TΔS. Endothermic reactions with positive ΔS become spontaneous at high temperatures; exothermic reactions with negative ΔS become spontaneous at low temperatures.
It's the temperature where ΔG = 0, calculated as T = ΔH/ΔS. Below this temperature, one sign of ΔG dominates; above it, the other does.
ΔG° = -RT ln K. A negative ΔG° means K > 1 (products favored). Every ~5.7 kJ/mol of ΔG° corresponds to about one order of magnitude change in K at 298 K.
ΔG° is the free energy change under standard conditions (1 bar, 1 M). ΔG is the actual free energy change: ΔG = ΔG° + RT ln Q, where Q is the reaction quotient.