**Free Energy**

Rolling a ball downhill is much easier than rolling a ball uphill. The same is true of chemical reactions. Reactions that produce or release energy occur much more easily (spontaneously) while those that take energy do not occur as easily (non-spontaneously).

- A
**spontaneous process**is one that, once started, proceeds on its own without any external influence. - A
**non-spontaneous process**is one that takes place only under the constant influence of an external influence.

Two key factors are at work in determining the spontaneity (Free Energy) of chemical reactions: enthalpy and entropy. This relationship is expressed using the following mathematical equation:

ΔG = ΔH - TΔS

where ΔG is the Free-energy Change (Gibbs Free Energy), ΔH is the enthalpy, and ΔS is the entropy. The remaining term, T, is for temperature. Recall from the module on Types of Matter, energy increases as temperature increases, which causes changes in the phase of matter. Since solids are much more ordered than liquids and gases, entropy will also change with the phase of matter, such that solids are low in entropy and gases are highest in entropy.

We can then tighten our definitions of sponataneous and non-spontaneous processes as:

- For a
**spontaneous process,**ΔG is negative, free-energy is released and is said to be**exergonic** - For a
**non-spontaneous process,**ΔG is positive, free-energy is added, and is said to be**endogonic.**

We can show the change in free-energy graphically using a reaction diagram, as shown below.

Scroll your mouse over the graphs to examine differences between exergonic and endogonic processes. | ||

Quiz yourself on your understanding of Energy