Against the Second Law of Thermodynamics
Daniel P. Sheehan, Beyond the Second Law of Thermodynamics: "These culminated in 2012-13 with a series of laboratory experiments that showed true second law breakdown. The demonstration was straightforward. A small, closed, high- temperature cavity contained two metal catalysts (rhenium and tungsten), which were known to dissociate molecular hydro- gen (H2) to different degrees (Figure 1). (Rhenium dissociates hydrogen molecules into atoms better than tungsten does; conversely, tungsten recombines hydrogen atoms back into hydrogen molecules better than rhenium.) Because the dissociation reaction (H2 -> 2H) is endothermic (absorbs heat), and the recombination reaction (2H -> H2) is exothermic (liberates heat), when hydrogen was introduced into the cavity, the rhenium surfaces cooled (up to more than 125 K) relative to the tungsten (Figure 2). Because the hydrogen-metal reactions were ongoing in the sealed cavity, the rhenium stayed cooler than the tungsten indefinitely. This permanent temperature difference--this steady-state nonequilibrium--is expressly forbidden by the second law, not just because the system won't settle down to a single-temperature equilibrium, but because this steady-state temperature difference can, in principle, be used to drive a heat engine (or produce electricity) solely by converting heat back into work, which is a violation of one of the most fundamental statements of the second law (Kelvin- Planck formulation)."
The second law of thermodynamics is almost obviously false for chemical systems. Consider the (valid) argument that, if catalysts can shift chemical equilibrium, the second law would be violated:
"In the presence of a catalyst, both the forward and reverse reaction rates will speed up equally, thereby allowing the system to reach equilibrium faster. However, it is very important to keep in mind that the addition of a catalyst has no effect whatsoever on the final equilibrium position of the reaction. It simply gets it there faster. (...) To reiterate, catalysts do not affect the equilibrium state of a reaction. In the presence of a catalyst, the same amounts of reactants and products will be present at equilibrium as there would be in the uncatalyzed reaction. To state this in chemical terms, catalysts affect the kinetics, but not the thermodynamics, of a reaction. If the addition of catalysts could possibly alter the equilibrium state of the reaction, this would violate the second rule of thermodynamics..."
It is evident that, for the dissociation-association reaction
A B C,
a catalyst cannot speed up both the forward and reverse reaction rates equally, due to the entirely different forward and reverse catalytic mechanisms. In the forward (dissociation) reaction, the catalyst should just meet and split A. In the reverse (association) reaction, the catalyst should first get together B and C, which, if the diffusion factor is predominant, could be highly improbable.
Catalysts do shift chemical equilibrium, in violation of the second law of thermodynamics.
I have started the same discussion (and it has developed in an interesting way) here:
Chemical Thermodynamics - Second Law / Entropy Review