18.3 Gibbs Free Energy and the Relationship between Delta G, Delta H, & Delta S | General Chemistry

Chad continues the chapter on Thermodynamics with an introduction to Gibbs Free Energy. Simply put, Gibbs Free Energy is the energy available to do work. For a spontaneous reaction, the change in Gibbs Free Energy (Delta G) is negative. For a nonspontaneous process, the change in Gibbs Free Energy (Delta G) is positive. And for a reaction that has reach equilibrium, the change in Gibbs Free Energy (Delta G) is equal to zero.
Mathematically, Delta G is defined in terms of the enthalpy and entropy changes as follows:
Delta G = Delta H - T(Delta S)
A reaction for which Delta H is negative and Delta S is positive will be spontaneous at all temperatures.
A reaction for which Delta H is positive and Delta S is negative will be nonspontaneous at all temperatures (or the reverse reaction will be spontaneous at all temperatures).
A reaction for which Delta H is negative and Delta S is negative will be spontaneous at low temperatures.
A reaction for which Delta H is positive and Delta S is positive will be spontaneous at high temperatures.
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00:00 Lesson Introduction
00:24 Introduction to Gibbs Free Energy
04:13 Derivation of Gibbs Free Energy
09:57 Delta G = Delta H - T(Delta S)
14:25 Negative Delta H, Positive Delta S
15:51 Positive Delta H, Negative Delta S
17:11 Positive Delta H, Positive Delta S
22:16 Negative Delta H, Negative Delta S
26:25 Delta H, Delta S, and Temperature of Phase Changes
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So at 26:08 you have the Haber process on the board which I heard happens at like 400-450C but if you calculate T = delta H / delta S for the point where the reaction stops occuring spontaneously I got 463.64 K or 190.49 C, which is less than the actual temp at which its carried out industrially, this confused me at first but apparently its bc of le chaliter's principle -- its carried out at high enough pressure to favour the side with low moles. I just thought i would comment this if anyone had the same Q or something

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