Quote: Originally posted by SplendidAcylation |
It appears you have fallen foul of DraconicAcid, he loves to correct the difference between kinetics and thermodynamics, he has done it to me twice!
That was a few years ago but the difference between the two is still fuzzy to me sadly.
"WoW big concepts, very small head." - I can relate
| Well, they are completely different. Try to think of it this way: thermodynamics can answer the question
“will I get X product?” But it won’t tell you whether getting to that product will take a second or (literally) a trillion years. That’s where
kinetics comes in.
This is where it’s helpful to look at a free energy diagram (shamelessly stolen from some website):
So looking at this diagram, you can see that the ΔG0 is the difference in energy between the reactants and products. In this case, it’s
negative, so the reaction is spontaneous. Great. However, there’s also the ΔG‡, otherwise known as Ea, or activation energy. This is
what tells you how fast the reaction is (kinetics!). A higher Ea means a slower reaction. A lower Ea indicates a fast reaction. Generally we supply
this energy by simply heating the reaction mixture, which is why (you may have been wondering) many “spontaneous” reactions still need to be
heated. Sometimes the Ea is too high to be overcome just by applying heat, though. It may be that your starting materials or products undergo
decomposition or side reactions at the temperatures required. In these cases, the reaction won’t work from a practical standpoint even if it is
thermodynamically favorable, but it may be possible to lower the activation energy by adding a catalyst. That’s what catalysts do: they provide an
alternative reaction pathway with a lower activation energy, thereby making the reaction proceed faster.
I like free energy diagrams because they combine thermodynamics and kinetics into one nice visual. |