Final answer:
The melting of ice at 0 °C has a positive ΔH and ΔS with ΔG of zero at equilibrium. Below 0 °C, ΔG would be positive, indicating non-spontaneous melting, but this condition is not reflected in the provided options.
Step-by-step explanation:
The melting of ice below 0 °C at 1 atm pressure is not spontaneous, as the equilibrium melting point is 0 °C. When considering the statements given, we focus on the melting at 0 °C, where the process is spontaneous. The enthalpy change (ΔH) for the melting of ice is positive because heat is absorbed by the ice, which is an endothermic process. The entropy change (ΔS) is also positive because the melting of ice results in an increase in the disorder of the water molecules. The Gibbs free energy (ΔG) of the system at 0 °C and equilibrium is zero, indicating a spontaneous process at this specific temperature.
However, since the question asks for melting below 0 °C, recall that at temperatures below the melting point, the system is not at equilibrium and the process is not spontaneous. Thus, the Gibbs free energy (ΔG) would be positive, indicating a non-spontaneous process in this condition.
Given these facts, the statement that correctly describes the melting of ice at 0 °C (which is the standard condition for discussing phase changes) is 'ΔH is positive; ΔS is positive; ΔG is negative' for the spontaneous process. However, since the meling below 0 °C is non-spontaneous, ΔG would actually be positive, which is not listed in the options provided. Therefore, there seems to be a misunderstanding in the question as none of the listed options accurately reflect the thermodynamic quantities for melting ice below 0 °C.