How this instrument works
Gibbs free energy change, ΔG, is chemistry's single verdict on whether a reaction happens spontaneously at a given temperature and pressure, without needing an outside energy source to force it along. A negative ΔG means the reaction is spontaneous as written — it will proceed on its own, though not necessarily quickly. A positive ΔG means the reverse reaction is the spontaneous one instead. A ΔG of exactly zero means the reaction sits at equilibrium, with forward and reverse rates perfectly balanced.
ΔG combines two competing thermodynamic drives into one number: enthalpy change, ΔH, which measures whether a reaction releases heat (negative ΔH, exothermic) or absorbs it (positive ΔH, endothermic), and entropy change, ΔS, which measures whether a reaction increases disorder (positive ΔS, favoured) or decreases it (negative ΔS, disfavoured). The formula ΔG = ΔH − TΔS weighs entropy's contribution by temperature — TΔS — so entropy matters more at high temperature and less at low temperature, while enthalpy's pull stays the same regardless of temperature.
This is exactly why some reactions are spontaneous only above or only below a certain temperature: a reaction with both ΔH and ΔS negative (releases heat but decreases disorder) can be spontaneous at low temperature, where the enthalpy term dominates, and non-spontaneous at high temperature, where the TΔS term grows large enough to flip the sign of ΔG. Reading off exactly where that crossover happens is one of the most practical things ΔG's temperature dependence is used for.
- Enter the reaction's enthalpy change into Enthalpy change, delta H (kJ/mol) — negative for exothermic, positive for endothermic.
- Enter the absolute temperature into Temperature (K) — kelvin only, and it must be greater than zero.
- Enter the reaction's entropy change into Entropy change, delta S (J/(mol*K)) — note the units are joules, not kilojoules, per mole per kelvin.
- Read the verdict off Gibbs free energy change, delta G (kJ/mol): negative means spontaneous as written, positive means non-spontaneous, zero means equilibrium.
- Watch the unit mismatch built into the formula — ΔH and ΔG are in kJ/mol while ΔS is in J/(mol·K); the instrument divides the TΔS term by 1000 internally so you don't have to convert by hand.
Worked example — ammonia synthesis (the Haber process)
Enter −92.4 into Enthalpy change, delta H (kJ/mol), 298 into Temperature (K), and −198.7 into Entropy change, delta S (J/(mol*K)) — standard textbook values for N2(g) + 3H2(g) → 2NH3(g) at room temperature. Gibbs free energy change, delta G (kJ/mol) reads −33.1874 kJ/mol.
By hand: ΔG = −92.4 − (298 × (−198.7) / 1000) = −92.4 − (−59.2126) = −92.4 + 59.2126 = −33.1874 kJ/mol. The negative result confirms ammonia synthesis is spontaneous at room temperature, even though the reaction's entropy drops sharply (4 moles of gas become 2, so the system becomes more ordered) — the strongly exothermic ΔH more than compensates for that unfavourable entropy term at 298 K.
Questions
What does a negative ΔG actually tell me?
It tells you the reaction, as written, is thermodynamically spontaneous at the temperature you entered — it will proceed on its own toward products without needing continuous outside energy input. It says nothing about speed: a reaction can have a strongly negative ΔG and still take years to happen if its activation energy is high, since ΔG describes whether a reaction is favoured, not how fast it gets there.
Why does temperature change whether a reaction is spontaneous?
Because ΔG weighs entropy's contribution by temperature through the TΔS term, while enthalpy's contribution, ΔH, stays fixed regardless of temperature. A reaction with a negative ΔH but also a negative ΔS can be spontaneous at low temperature, where TΔS is small and ΔH dominates, and become non-spontaneous at high temperature, where TΔS grows large enough to outweigh ΔH and flip ΔG positive.
Why are ΔH and ΔS given in different units?
Purely by convention — enthalpy changes for typical reactions run in the tens to hundreds of kJ/mol, while entropy changes run in tens to hundreds of J/(mol·K), so chemists report each in whichever unit keeps the numbers in a convenient range. This instrument's formula divides the ΔS term by 1000 internally to convert it to kJ before combining it with ΔH, so you can enter each value in its usual textbook unit directly.
What does ΔG = 0 mean?
It means the reaction sits exactly at equilibrium at that temperature — the forward and reverse reactions proceed at exactly matched rates, so there's no net drive pushing the reaction toward either products or reactants. Solving ΔH = TΔS for T (rearranging ΔG = 0 = ΔH − TΔS) gives the specific crossover temperature at which a reaction switches from spontaneous to non-spontaneous, or vice versa.
Can a reaction be spontaneous even if it absorbs heat?
Yes — an endothermic reaction (positive ΔH) can still have a negative ΔG, and therefore be spontaneous, if its entropy increase is large enough that the TΔS term outweighs the positive ΔH, especially at high temperature. Ice melting above 0°C is a familiar everyday example: it absorbs heat (positive ΔH) but the large entropy increase of turning ordered solid into disordered liquid makes ΔG negative at those temperatures.