If we have solid water (ice) at -5 C and atm P, then according to thermodynamics, the process that takes that ice (at -5 C & atm P) from solid to liquid is non-spontaneous, which means its ΔG > 0. Hence, it is impossible for someone to observe ice melt to liquid at -5 C, and assuming we started with a solid block of ice, this means there is no liquid water in the system. However, if I have an arbitrary chemical reaction that is happening, with a ΔG > 0 (endergonic), there will always be an equilibrium established (given enough time) that will have some reactants AND some products present (the amount of each present depends on the K(eq) value or equivalently the ΔG). My misunderstanding comes from why these two cases are different? Why, in the case of the phase transition, is it thermodynamically impossible to have both reactants and products present (ice and liquid), whereas in the chemical reaction, even if ΔG > 0 like before, its expected that both reactants and products will be present in the system? Hopefully everything here makes sense.