One mole of H2O initially exists as ice at –10 °C. The substance is brought to 100 °C in several distinct steps:
Step 1: Warming the ice from –10 °C to 0 °C.
Step 2: Melting the ice at 0 °C.
Step 3: Warming the liquid water from 0 °C to 100 °C.
Step 4: Vaporizing the water at 100 °C.
For water, ΔHfus = 6.01 kJ mol−1 at 0 °C and ΔHvap = 40.7 kJ mol−1 at 100 °C. Assume the molar heat capacity of ice is 37.5 J mol−1 K−1 and that of liquid water is 75.3 J mol−1 K−1, both taken as temperature-independent over the ranges used.
(a) Using the information provided, calculate ΔS per mole for Step 2 and for Step 4. Report each value with units.
(b) Without detailed calculations, qualitatively compare the entropy changes for Steps 1 and 3 with those for Steps 2 and 4. Explain your reasoning in terms of matter and energy dispersal.
(c) Identify which single step contributes the largest fraction of the total entropy increase for the overall transformation from ice at –10 °C to steam at 100 °C. Justify your answer using the results of part (a) and the qualitative reasoning from part (b).