Final Answer:
The non-spontaneity of a reaction refers to its overall thermodynamic feasibility, while the forward direction proceeding indicates a kinetic possibility. In this case, although the overall reaction is non-spontaneous (ΔG∘ = 4.76 kJ/mol), the forward direction can still occur under certain conditions, allowing the system to reach equilibrium with a small concentration of products.
Step-by-step explanation:
The non-spontaneity of the reaction, as indicated by the positive ΔG∘ value, suggests that the system would not spontaneously move towards the formation of products under standard conditions. However, thermodynamics does not provide information about the rate at which a reaction occurs. The kinetics of a reaction consider the speed at which reactants transform into products, and in this scenario, the forward reaction may proceed slowly but still reach equilibrium.
To understand why the reaction proceeds in the forward direction, consider the relationship between Gibbs free energy (ΔG), enthalpy (ΔH), and entropy (ΔS) in the Gibbs-Helmholtz equation: ΔG = ΔH - TΔS. Even if ΔG∘ is positive, the temperature-dependent term (TΔS) can make the overall ΔG negative under certain conditions. Hence, at a specific temperature, the entropy-driven term can overcome the enthalpic barrier, allowing the forward reaction to occur.
In summary, while the reaction is thermodynamically unfavorable in terms of ΔG∘, the kinetics of the reaction, influenced by temperature and entropy changes, enable the forward direction to proceed, albeit slowly, and reach equilibrium with a finite concentration of products.