Second Law And Entropy

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Melting of ice at 273 K and 1 atm is a spontaneous process. Which of the following correctly describes the thermodynamic parameters at the melting point?

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1. **Nature of Melting**: Ice → Water is endothermic (ΔH > 0) because energy is needed to overcome hydrogen bonds in the crystal lattice. 2. **Entropy Change**: The liquid state is more disordered than the solid, so ΔS > 0. 3. **At the Melting Point**: The solid and liquid phases coexist in equilibrium at 273 K and 1 atm. At equilibrium, $\Delta G = 0$ (neither phase is thermodynamically favoured over the other). 4. **Consistency Check with Gibbs Equation**: $\Delta G = \Delta H - T\Delta S = 0 \Rightarrow T_{mp} = \frac{\Delta H}{\Delta S}$. Both ΔH and ΔS are positive, confirming T_mp is a positive, finite value. 5. **Why Not the Others?** - ΔH < 0 would describe an exothermic transition (e.g., freezing, not melting). - ΔS < 0 would imply ordering, which contradicts the solid-to-liquid transition. - ΔG > 0 would mean the process is non-spontaneous, contradicting equilibrium at the exact melting point.

Second Law And Entropy

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