Looking at Cells Under the Energy Microscope

When we discussed heat and entropy, it became crystal clear that thermodynamics is critically important to cells and cell function. Now, let's take a closer look at the involvement of thermodynamics in cellular biochemical reactions. The second law of thermodynamics (yep, it's back for more) can be interpreted in the context of chemical reactions, where energy transfer events tend to progress downhill—that is, products have less energy than reactants—and any extra energy produced during chemical reactions is lost as heat, which ultimately increases the disorder in the universe. In other words, once an energy transfer event occurs, there is less energy left in the product to do additional work.

Cellular reactions that occur spontaneously will proceed to a more disordered, but lower energy state. This may seem counterintuitive if you were thinking that increasing disorder always means increasing energy in the products, but this is not the case. Remember that cellular reactions release heat to the surroundings as they form products that are themselves in a lower energy state.

Proteins that have been broken down into their individual amino acids (see the Biomolecules and the Chemistry of Life unit if this is not ringing a bell) by a cellular reaction will not suddenly spontaneously re-form intact proteins. Instead, any uphill reaction is a reaction that involves a large input of energy, such as the reactions that build proteins or nucleic acids. However, when we start trying to predict if a reaction will occur, or why a protein takes on a certain structure, it is not enough to rely on the second law of thermodynamics. Instead, we need to combine the first and the second laws of thermodynamics.

The first and second laws of thermodynamics were combined into one equation by Josiah Willard Gibbs (not Barry, Robin, or Maurice) in the late 1800s. The equation is

ΔH = ΔG + TΔS,

or the change in something called enthalpy (ΔH) equals the change in something called free energy (ΔG) plus the absolute temperature (T, in degrees Kelvin, where x degrees Kelvin = y degrees Celsius + 273) multiplied by the change in entropy (ΔS). Did you catch all that?

This equation is useful (it is; stop shaking your head) because it allows scientists to find out information about a reaction while only knowing a few details about the system. This equation is not nearly as equation-y as it may seem at first glance.

You might say, "But Shmoop, what exactly is enthalpy?" Sorry, you are not privileged enough to receive such information. In all seriousness, you will learn more about enthalpy in chemistry, but for our purposes, enthalpy is equal to the total amount of energy in a system.

If a system gains heat from a chemical reaction, ΔH will be positive. Alternatively, if heat is lost from a system, ΔH will be negative. Not hard, right? Then, there is the term TΔS. We know that entropy is a measurement of disorder. In this case, TΔS takes into account only the entropy change of the system.

• If a reaction increases the disorder in a system, the entropy term TΔS will be positive.
   
• If a reaction decreases the disorder in a system, the entropy term will be negative.

Now, what about free energy? You wish gasoline were free... Oh, right, ΔG. Coming right up.

Brain Snack

Spontaneous combustion—when something starts to be consumed by fire without external ignition—is real! And, in reference to our enthalpy discussion above, spontaneous combustion occurs if the heat in the system increases without being able to escape.