Final answer:
The reaction Y -> X is initially spontaneous with a negative ΔG. However, an excess of product X would shift the equilibrium and make the reverse reaction X -> Y become favored and spontaneous. This phenomenon follows Le Chatelier's principle regarding changes in a system at equilibrium.
Step-by-step explanation:
Although the reaction Y -> X is spontaneous due to the lower energy of X, a significant excess of X would make the reverse reaction favoured with a negative ΔG (delta G). In chemistry, a spontaneous reaction is one that releases free energy, indicating that the sign of ΔG must be negative, which means the products have less free energy than the reactants. This scenario can be altered by changes in concentration, as increasing the concentration of the product X beyond a certain point would make the reverse reaction, X -> Y, more favourable in terms of having a negative ΔG.
For a reaction that is exothermic and has a negative change in free energy, it is termed exergonic. However, adding more of the product X would shift the equilibrium towards the reactants, favouring the reverse reaction which, under these new conditions, would have a negative ΔG thus becoming spontaneous in the opposite direction. This principle follows Le Chatelier's principle where a system at equilibrium will try to counteract any changes imposed on it such as changes in concentration.