On Delta-G

The quantity know in thermodynamics as ΔG expresses the free energy variation that occurs with a chemical reaction or phase transition (and other processes). The letter G was chosen as na homage to the American scientist J.W. Gibbs, who did work of fundamental importance for themodynamics.

Restrictin our scope to chemical reactions, Gibbs free energy variation consists of two terms:

ΔG = ΔH - TΔS

ΔH is the enthalpy variation - in other words, the amount of heat produced or absorbed by the reaction; or the difference in heat content of products and reactants.

ΔS is the entropy variation. Entropy is a fascinating but often misunderstood quantity; it can be interpreted as a measure of disorder in a system; ΔS is the difference between entropy of products and reagents. Higher entropy corrsponds to more disorder.

ΔG is negative for spontaneous reactions. Because in nature everything evolves towards the minimum energy state.

Considering the equation above, it can be seen that ΔH and ΔS do not need to be both favorable; one of them can be unfavorable but if the other is big enough, the reaction will still be spontaneous.

ΔH is the heat of reaction; it can be determined experimentally or estimated using appropriate data tables (and it is already known for many reactions). There isn't much to say about this quantity, though it is worth remembering that stable products (water, carbon dioxide, elemental nitrogen) have low enthalpy. Hence, reactions with these products are often esothermic.

ΔS instead is open to more discussion. Natural systems also evolve towards higher entropy states, so a positive ΔS is favorable. There are rules of thumb regarding entropy variation: remember, it measures disorder. Solid reagents have very low entropy (because entropy is also a function of spatial distribution of molecules); liquids have more and gases have very high entropy. Breaking up big molecules into smaller ones causes in entropy increase (combustion, for example). Formation of small, volatile molecules (polycondensation liberating water, for example) often increases entropy, while joining together small molecules into bigger ones (chain polymerization) decreases it. Isomerization from an ordered structure to a less ordered one produces an entropy increase, and viceversa.

And today's lesson is over, ok?