[poobear]

Hi!

In a spontaneous reaction the enthalpy normally decreases. As such the temperature of the surrounding normally goes up. So the enthalpy of the entire system does not change. What only happens is that the enthalpy gets less localized, from just the substrate to the entire system.

This is my view of a chemical reaction, please correct me if I'm wrong.

Can I therefore view enthalpy as an "energy entropy"? Because in a spontaneous reaction it normally gets less localized.

I was thinking about Gibbs energy (∆G = ∆H - T∆S), and was wondering if H and S should be looked as the energy entropy and state entropy, respectively. I'm starting to think that everything is entropy in different forms (i.e. delocalization of the form), but I'm not sure if that's correct.

Thanks!

[Yggdrasil]

The relationship ΔG <= 0 (i.e. that reactions seek to minimize their free energy) is derived from the second law of thermodynamics (specifically for the case of a reaction occurring at constant pressure).  The second law states that the entropy of the universe is always increasing.  There are two components to the entropy of the universe which we must consider, the entropy of the system and the entropy of the surroundings.  By splitting the entropy of the universe into these two components, we can write the second law as follows:

ΔS_uni = ΔS_sys + ΔS_surr >= 0

From this equation, we can see that there will be trade offs between the entropy of the system and the entropy of the surroundings.  In order for the system to lose entropy and become more ordered, the surroundings must gain entropy to compensate and vice versa.

The formula for the Gibbs free energy reflects these trade offs between the entropy of the system and the entropy of the surroundings.  Recall that for a reaction occurring at constant pressure, the change in enthalpy is equal to the amount of heat released/absorbed by the system, ΔH = q.  Further, remember that the change in enthalpy of the surroundings is given by the following formula: ΔS_surr = q/T_surr.  Therefore, the quantity ΔH/T_surr will tell you about the change in entropy of the surroundings. 

Given that ΔS_surr = -ΔH/T_surr, we can plug this value into our expression for the entropy of the universe to give:

ΔS_uni = ΔS_sys + ΔS_surr =  ΔS_sys - ΔH/T_surr

Applying the second law, we obtain the inequality:

ΔS - ΔH/T >= 0

Or, equivalently:

ΔH - TΔS <= 0

Which is exactly the equation that explains why spontaneous reactions tend to minimize their free energy.


From these derivations we can see that exothermic reactions are favorable because they transfer heat to the surroundings and therefore increase the entropy of the surroundings.  As you correctly mentioned, you can think of this heat transfer as increasing entropy because you "delocalize" the energy trapped within the system to occupy the much larger surroundings.

[poobear]

Great answer <3