Seven Perspectives on Free Energy

Ah, free energy: the eternally confusing concept that lies at the heart of thermodynamics and statistical mechanics.

After so many years, I still find it remarkably difficult to properly interepret the free energy of a system, because there's so many different ways to define it and think about it. Each time you encounter the free energy, you have to cope with a different and often unrelated viewpoint. And then at the end of the day, once you've collected a bunch of diverse grab-bag perspectives, you have to somehow consolidate and reconcile them together into a coherent understanding.

Sometimes I wish there were one definition of The Free Energy, but the more and more I read about it, the less and less I believe in any unified understanding. At its core, free energy is a multi-faced beast.

Typically, we first encounter this confusing creature in high school chemistry class, when we learn about what makes reactions ‘‘go’’ – but as we continue through our educational journey, the unweildy animal pops its head at us again and again in unexpected places – in the biophysics of molecular motors, in the thermodynamics of steam engines, in discussions of protein folding, in the modern statistical mechanical theory, here, there, everywhere. What lies in common among all these definitions of free energy? What understanding can we glean from each of the interpretations? And how do they help build towards a cohesive understanding of free energy?

I don't have an answer to these questions. Nevertheless, I find it enlightening to bring together multiple perspectives towards free energy in one same place – mostly to soothe my own cognitive dissonance towards the whole confusing affair.

Here, I present to you: Seven Perspectives on Free Energy.


We begin our journey with a qualitative explanation from high school chemistry class: the free energy tells you whether matter wants to transform. Leave iron in the rain and it becomes rust; put water in a freezer and it becomes ice; put water in an oven and it becomes steam. In all these chemical reactions, matter begins in one form and ends in another, and the driving force behind all such transformations is a decrease in free energy.

We can figure out whether chemical reactions proceed by comparing the free energy of their reactants and products. If a reaction results in a free energy that's smaller than what you started out with, then the transformation occurs. Typically, we define Delta F = F_{products} - F_{reactants} to quantify how much the free energy changes due to a chemical reaction; if it's negative (Delta F < 0) then we know the reaction occurs spontaneously because F_{products} < F_{reactants}. Conversely, if the free energy increases over the course of a reaction (Delta F > 0), then it can't happen spontaneously.

This simple rule reveals some interesting physics:

The timescale question

Notice that our description of chemical reactions also has no mention of how long it takes for the transformation to occur. Indeed, the biggest caveat of thermodynamics is that it assumes you've waited an infinite amount of time for things to happen!

In reality, just because Delta F < 0 doesn't necessarily mean that a transformation will occur in experiments. All it says is that if you wait for infinitely long, then the transformation will ventually occur – but it doesn't tell you whether you have to wait milliseconds, or hours, or millenia, or the lifetime of the universe. Often the system must overcome an activation energy barrier to undergo a reaction; if it is too costly to ‘‘go over the hump’’ then the reaction might not actually occur over timescales we care about.

Here are a few instances where kinetics hinder a thermodynamically favored reaction (on human timescales):

In summary, the statement that matter transforms into a state that minimizes its free energy is deceptively subtle. It is a simple rule, and yet it includes many assumptions and implications. We must keep these assumptions in mind whenever we do thermodynamics.


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