Gibbs energy of reaction
The Gibbs energy of reaction, written ΔG, is the single number that tells you which way a chemical reaction wants to go, and how strongly. Picture a hill between reactants and products: ΔG is the difference in 'free-energy height' between the two sides. If products sit lower, ΔG is negative and the reaction rolls forward on its own; if they sit higher, ΔG is positive and the reaction prefers to run backward.
Formally, ΔG is the change in the system's Gibbs free energy as a reaction proceeds, equal to the free energy of the products minus that of the reactants under the current conditions. The standard value, ΔG°, is the same comparison taken with everything in its standard state. The two are linked by how far the mixture is from equilibrium, captured through the reaction quotient.
Why it matters: ΔG is the working tool of chemical thermodynamics — it predicts feasibility, sets the maximum work and voltage a reaction can supply, and pins down the equilibrium constant. The honest caution: a negative ΔG only certifies that a reaction is thermodynamically allowed. Whether it actually proceeds at a useful rate is a separate matter decided by kinetics, which ΔG knows nothing about.
Burning glucose has a strongly negative ΔG of about −2,880 kJ per mole, so it's hugely favorable. Yet a sugar cube sits unchanged on the table for years — favorable, but with no fast path until heat or an enzyme opens one.
A very negative ΔG means very favorable — but says nothing about speed.
Don't confuse ΔG with ΔG°. ΔG° is fixed for a reaction (a tabulated constant at a given temperature), while ΔG shifts as the reaction proceeds and the concentrations change — reaching zero precisely at equilibrium.