Acids and Bases

Buffers

Redox

Chirality

Entropy


Home



Entropy

Entropy is a great source of confusion among biochemistry students, especially regarding how the concept is applied to explaining the chemistry and processes that take place in living things. A common definition for entropy is:

Entropy (S) is the measure of disorder for a system.

This brings us to the second law of thermodynamics:

The entropy of an isolated system increases for a spontaneous change.

Try to think of entropy as a measure of the possible states of a system. For example, if you were to ask yourself which has more entropy, ice or liquid water, you could reason that the molecules in the liquid state are relatively free to move and rotate, whereas in ice the molecules are "frozen" in fixed positions, with little movement except for the vibration of the bonds. So for water, there are much more orientations possible for each molecule, resulting in a multitude of possible states for the liquid solutions.

Another example to contemplate is the mixing of two liquids such as acetic acid and water. If acetic acid is poured into water, the acetic acid molecules will spontaneously disperse themselves evenly throughout the solution, and will not stay bunched together and separated from the water. This is because, in the case where the two remain neatly separated, the situation is much more ordered (less random), and the natural tendency is for randomness (entropy) to increase.

"Aha! But what about oil and water?", you might ask. The process of mixing oil and water does not disobey the second law of thermodynamics, for it is a more complicated scenario than acetic acid and water. Oil is a hydrophobic substance that does not H bond to water, yet because it occupies space in an aqueous solution, the water molecules must rearrange their hydrogen bonds to form a clathrate structure that surrounds the nonpolar oil like a cage. This clathrate structure introduces a lot of order to the system, for the water molecules at the oil-water interface are much more restricted in their movement. Because the smaller the oil-water interface the less order is introduced into this system, the system will equilibrate to minimize the surface area of the oil-water interface. The best way to do this is to keep all of the oil together in one glob, so that the volume-to-surface-area ratio is maximized. If the oil dispersed itself similarly to the acetic acid discussed previously, there would be many more water molecules involved in clathrate structures. So this solution sacrifices the entropy increase associated with even dispersal of a solute, for the benefit of a reduced amount of order introduced by a clathrate structure.

A more biochemical example is the folding and unfolding of proteins. A polypeptide that has folded into a functional protein is much more restricted in orientation than an unfolded polypeptide, which is like a string that can be bent and twisted into countless conformations. So for a protein that undergoes folding, the entropy change is negative, whereas for a protein that unfolds, the entropy increases, and the change is positive.

Chemists use entropy as a tool for understanding spontaneous change. For a reaction, entropy can be used to predict whether it will proceed spontaneously. However, be very careful when attempting to use entropy as a tool for detecting spontaneous change in biochemistry, for you must remember that this only applies to closed (isolated) systems, thus not including biological systems. Organisms are open to the universe because they respire, excrete waste, generate heat, etc. So the only way to use entropy as a determinant for change in an open system is if you consider the entropic effects on the entire universe, which is very difficult to measure.

Because the organism is not a closed system and because of the difficulties of measuring changes in entropy for biological processes, the Gibbs Free Energy (G) is a much better indicator of spontaneous change, for it takes the enthalpic energy changes into account as well.

DG = DH - TDS

For spontaneous processes, the DG, or change in free energy, is negative (-). It is this DG that provides a very useful gauge as to whether a reaction will proceed in either the forward or backward direction in a given environment and with a specific concentration of reactants.

To give you an idea of why the free energy is a better way to determine whether a reaction is spontaneous or not, note that while protein folding leads to a decrease in entropy, it is still a spontaneous process inside of cells. Why is this? The enthalpic contributions of the intrachain hydrophobic packing, hydrogen bonding, ionic interactions, and disulfide bonding counteract the entropy decrease upon folding. You see, these bonds and interactions release energy as they are formed, most of which is dissipated as heat that warms the cell, the organism, and the air surrounding the organism. This heat does contribute to the increase of entropy of the universe, for it makes molecules move about and rotate faster, leading to more disorder. Yet the entropy due to heat vented to the universe is extremely difficult to quantitate for an isolated, complex, biochemical process in an organism, while enthalpy changes from the process (such as bond formation and the interactions between molecules) are more easily evaluated by experimentation.



Copyright © 1999, Harcourt College Publishers, A Harcourt Higher Learning Company.
All rights reserved. Use of this site indicates acceptance of the Terms and Conditions of Use.
For problems or suggestions concerning this service, please contact Chemistry Webmaster.