Just a bit of Thermodynamics
In module 7 was discussed a piece of thermodynamics of reactions in general, as well as the concept 'spontaneous'.
Imagine the following chemical system:
N2 + 3H2
2NH3 (ΔH < 0 and ΔS < 0)
Such a chemical system means for example that Nitrogen, Hydrogen and ammonia are toghether, reacting, in a cylinder under a piston.
Now remember that formula:
ΔG = ΔH - TΔS
The forward reaction (production of ammonia) is exothermic (ΔH < 0)
and that garantees a certain spontaneity of the reaction to the right.
On the other side however: the backward reaction inplicates an increase of entropy (the number op particles increases, doubles in the example, and thus, the amount of chaos increases when reacting from right to left.
This also garantees a certain spontaneity for the backreaction. (ΔS increases in the reaction backwards and decreases in the reaction to the right).
So in the example both reactions, the one to the right and the one back, have a certain spontaneity for different reasons.
In the formula ΔG = ΔH - TΔS, applied for the forward reaction, ΔH as well as ΔS have a negative value.
Mathematically this means that ΔH makes the value of ΔG more negative, where -TΔS will make the value of ΔG more positive.
It may happen that the two (ΔH= TΔS) have equal values, dependent on the circumstances.
Now we touch a thermodynamic secret:
You see: for understanding well the chemical equilibrium, you must know the concept of 'entropy'.
- At the moment that the value of ΔG equals ), chemical equilibrium is reached.
- When ΔG > 0, meaning that the free energy of the system increases, the backreaction will dominate (the equilibrium dislocates to the left)
- When ΔG < 0, meaning that the free energy of the system decreases, the forward reaction will dominate (the equilibrium dislocates to the right).