free energy and spontaneous reactions

Picture a ball on a hillside. Left alone, it rolls downhill on its own; to get it back up, you have to push and spend effort. Chemical reactions behave the same way: some run downhill all by themselves and even give off usable energy, while others only go if you push energy in. Free energy is the bookkeeping that tells you which is which.

Free energy (often written G, for Gibbs free energy) is the portion of a reaction's energy that is actually available to do useful work, after accounting for the disorder that thermodynamics demands. What matters is the change in free energy, written as a difference. If a reaction releases free energy, that change is negative, and the reaction is called spontaneous, or exergonic — it can proceed on its own and can power other tasks, like the downhill ball. If a reaction requires free energy, the change is positive, and the reaction is non-spontaneous, or endergonic — it needs an energy push, like rolling the ball uphill. Importantly, spontaneous does not mean fast; it only means the reaction is energetically downhill.

Cells live by this accounting. Breaking down food is downhill and releases free energy; building complex molecules is uphill and demands it. The cell's trick is to link the two so the energy from a downhill reaction pays for an uphill one. Free energy is therefore the single concept that lets us predict whether a cellular reaction can happen at all and whether it can be made to drive something useful.