Entropy describes the degree of disorder or “randomness” in a substance. For example a low temperature solid material has relatively low entropy compared to the material in gaseous form at elevated temperatures.
In the gaseous phase, atoms travel in any direction, occupying different regions in space, colliding together and with their container. The solid, a much more ordered system, has atoms neatly arranged in fixed positions, vibrating gently but not swapping places.
A Model for Entropy and Free Energy
Imagine a classroom of children, each sitting at desks in rows, quietly concentrating! This illustrates a fairly ordered “low entropy” system. Minutes later, the children leave the classroom and play outside. The running, shouting children, free to move in any direction, present a more disordered “high entropy” picture.
This analogy illustrates another aspect of low and high entropy systems. The children are now required to do some “work”, requiring great concentration. From the teacher’s perspective, which collection of youngsters would have the greater potential to complete this work effectively and efficiently, in the shortest time?
Naturally, the “low entropy” children complete the work required more successfully than the students running randomly, helter-skelter outside. Children in serried ranks in the classroom are quickly and effectively marshalled to complete the work required. Doing the same for the “high entropy” children outside would prove highly problematical!
Similarly once a chemical system has entered its high entropy phase, its capacity to do useful work (called “free energy”) diminishes considerably, whereas highly ordered, low entropy systems can be effectively harnessed to do work. The transfer of heat and its effective use is the realm of thermodynamics.
A “Real Life” Example of Chemical Entropy at Work
A familiar, chemical example, of entropy illustrates these points. Imagine some charcoal sitting below a juicy barbeque steak . Carbon atoms within its cold solid structure have a high degree of order (low entropy). The “work” to be extracted from charcoal is heat energy that will cook the steak to perfection. The “potential” of the charcoal to do useful work is very high.
To extract this work/heat energy, the charcoal must burn in oxygen from the surrounding air. The barbeque is lit! Once burning, the charcoal releases heat energy, completes the “work” in hand and the steak is beautifully cooked. The product of burning carbon in oxygen, however, is hot carbon dioxide. This highly disordered, rapidly dispersing, collection of molecules spreads out into the surrounding air carrying their “ability to work” (more heat energy) with them.
To collect, organise and extract the heat energy from these carbon dioxide molecules is nigh on impossible, not least because their random, high entropy nature allows their heat energy to escape into the surroundings.
Those spotting the simplification in this argument (the entropy of the oxygen required for reaction has been overlooked) should note that the entropy of the hot carbon dioxide “products” is considerably higher that the “mean” entropy of the cold charcoal and the oxygen. Also, consider what happens to the randomness of the surrounding air molecules during the reaction.
Combustion is an exothermic reaction, releasing energy to the surroundings. This energy, mainly heat, causes surrounding air molecules to speed up, moving more “randomly” than they would otherwise. In effect, the entropy of the surroundings (indeed the whole universe!) increases.
Professor Frank L. Lambert at Occidental College, Los Angeles has written a fairly extensive and entertaining guide to the fascinating subject of entropy, which is well worth a look.
