Talk:Introduction to entropy
The first paragraph
Feel free to knock the following down. I think the present (9 am PST, 26 October is good, but too dense, too many ideas per paragraph.) Maybe this is a better small step-wise start:
The concept of entropy is central to the second law of thermodynamics. A modern version of the second law is "Energy of all types spontaneously changes from being localized to becoming more dispersed or spread out, if it is not hindered". A simple example would be a hotter bar of iron touching a cooler bar. Always the energy of the hotter spreads out to the cooler until both are at the same temperature. Then, entropy (or better, entropy change, ΔS) is the quantitative measure of how much energy, q, has been dispersed in a process, divided by the temperature, T, at which it occurs. This yields the fundamental equation for entropy: q(reversible)/T.
(That "reversible" means that the process should be carried out at a temperature where each bar is almost the same -- the hotter only slightly warmer than the cooler so the energy would flow in reverse if there were only a slight change in the relative temperatures. When the temperature of one bar is considerably hotter than the other, as is most often the practical case, we can use calculus to 'paper-correct' the temperature difference. Our calculations simulate a series of reversible steps by treating the entropy change as though it occurred in tiny jumps or increments of temperature.)
There are two principal types of process involving energy change for which entropy measurement is especially valuable: heating (i.e., the transfer (spreading out) of energy from hotter to cooler and raising the temperature of 'the cooler'), or allowing the internal energy of a substance to become more spread out by allowing it to have more volume -- as in letting a gas expand or in mixing fluids (gases or liquids). Chemical reactions involve both types due to the release of energy (i.e., its dispersal to the surroundings) when new stronger bonds are formed as old ones are cleaved. (Endothermic reactions are the opposite: energy flows from the surroundings to become more dispersed in a system because that energy, in breaking bonds in the system, yield more particles that can more widely spread out that energy.)
The entropy of any substance at a given temperature, T, is given in tables of standard entropy as so many joules (energy)/T298 K. Actually, that final q/T 298 K is a sum and thus a ΔS of very many individiual measurements or calculations of q/T over very small ΔT changes from 0 K to T. Thus, the number for the entropy in joules/K is not the total amount of joules of internal energy in the substance (because the many increments of energy have each been divided by their T values). However, the standard entropy, listed in Tables at 298 K, is a useful rough comparison of the internal energy of every substance.EXAMPLE?? Graphite vs. diamond?? That is significant because it means the comparative value for the energy which a substance must have to exist at 298 K. (The phrase in italics is the meaning of "energy that is unable to do work" or "energy that is unavailable for thermodynamic work". Of course, the internal energy of a substance A can be transferred to another that is colder, but then substance A can no longer exist at its initial temperature.)