Thermodynamics

GIBBS FREE ENERGY

A thermodynamic indicator of success

We need to appreciate the principles of thermodynamics to understand how biochemical reactions occur. Section 1 explained how the first two laws of thermodynamics dictate the flow of energy through our universe. But how can we understand biochemical reactions in terms of these laws of energy flow? In the 1800s, a scientist named J. Willard Gibbs described a new value, “free energy,” to helps us to do exactly that. Gibbs Free Energy, denoted by G, describes the energy available to do work within a system, in this case a chemical reaction.

Free Energy can be used to make things occur that wouldn’t happen without a source of energy. Examples of work being done include the progress of a chemical reaction to create a new product, an engine running to turn the wheels of a car, and falling water moving the turbines of a hydroelectric power plant. Keep in mind that Free Energy, which is useful because it can be harnassed to do work, is different from the heat energy that is always lost in any process. This heat energy is generated as the cost of doing work, and is therefore often called nature’s “heat tax.” The Gibbs Free Energy Change, or DG, is the difference between the G of the reactants and the G of the products.

Thus, the Gibbs Free Energy of reaction is the maximum potential usable energy of a system, and is calculated by the following formula:

Energy diagram: The free energy profile of a reaction going from products to reactants.

Here is a reaction diagram that illustrates DG. If the DG is negative, meaning the products are at a lower energy than the reactants, then the reaction is thermodynamically favorable, meaning it can proceed. Remember, the second law of thermodynamics states that a system will always want to move toward a lower energy state. Thus, DG is useful because it tells us if a chemical reaction is likely to proceed in the direction it is written.

However, the DG does not tell us at all how fast this reaction will happen. This is because there is another quantity that is important in the energetics of a chemical reaction. Take another look at the energy diagram (above). Notice that as the reaction proceeds, the reactants actually have to move through a higher energy state before they can arrive at the low energy state of the products. In other words, the reactants have to scale an “energy wall,” called the Energy of Activation ( DG^‡). The higher the DG^‡, the taller the wall. As you might suspect, a taller wall is more difficult to scale, and thus reactions with a high DG^‡ proceed slowly. In contrast, reactions with a low DG^‡ progress more quickly. Such reactions are said to be kinetically favorable, or likely to proceed rapidly. (For more on biochemical kinetics, see the kinetics review.)

The figure above shows a reaction with a negative DG (the products have less free energy than the reactants). Such an energy-releasing reaction is said to be exergonic. However, some reactions have a positive DG (where products have more free energy than the starting reactants). As you might imagine, such reactions require energy input, and are called endergonic, or energy consuming.

Keep in mind that energy released in an exergonic reaction is not necessarily released as heat. Thus, an exergonic (energy releasing) reaction is not necessarily exothermic (heat releasing). In other words, a reaction with a negative DG (free energy) value may have either a positive or negative DH (enthalpy) value!