I was checking my answers for a chem test and came across a solution for a problem where the ∆S was pos, ∆H was neg and ∆G was neg. I believed it was only favorable at low temps, as reactions with a negative ∆H and negative ∆S are spontaneous at low temperatures. I was incorrect and the reaction was favorable at higher temps like 90 C as well. Is a reaction always favorable if the ∆G is negative (at all temps)? What about when the ∆S is positive? If it's not favorable in all situations, what temps would cause it to be favorable? What role does ∆H play in determining the favorability of a reaction?

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Explanation:

The second law of thermodynamics says that the entropy of the universe always increases for a spontaneous process: \Delta \text {S}_{\text{universe}}=\Delta \text {S}_{\text{system}} + \Delta \text {S}_{\text{surroundings}} > 0ΔS  

universe

​

=ΔS  

system

​

+ΔS  

surroundings

​

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At constant temperature and pressure, the change in Gibbs free energy is defined as \Delta \text G = \Delta \text H - \text{T}\Delta \text SΔG=ΔH−TΔSdelta, start text, G, end text, equals, delta, start text, H, end text, minus, start text, T, end text, delta, start text, S, end text.

When \Delta \text GΔGdelta, start text, G, end text is negative, a process will proceed spontaneously and is referred to as exergonic.

The spontaneity of a process can depend on the temperature.

Spontaneous processes

In chemistry, a spontaneous processes is one that occurs without the addition of external energy. A spontaneous process may take place quickly or slowly, because spontaneity is not related to kinetics or reaction rate. A classic example is the process of carbon in the form of a diamond turning into graphite, which can be written as the following reaction:

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