Energy in a body of water can be gained or lost depending on conditions. When water is heated above a certain temperature steam is generated. The increase in heat energy creates a higher level of disorder in the water molecules as they boil off and leave the liquid.
At the temperature at which a change of state occurs, the two states are in equilibrium with one another. For an ice-water system, equilibrium takes place at 0°C, so
is equal to 0 at that temperature. The heat of fusion of water is known to be equal to 6.01 kJ/mol, and so the Gibbs free energy equation can be solved for the entropy change that occurs during the melting of ice. The symbol
represents the entropy change during the melting process, while
is the freezing point of water.
The entropy change is positive as the solid state changes into the liquid state. If the transition went from the liquid to the solid state, the numerical value for
would be the same, but the sign would be reversed since we are going from a less ordered to a more ordered situation.
A similar calculation can be performed for the vaporization of liquid to gas. In this case we would use the molar heat of vaporization. This value would be 40.79 kJ/mol. The
would then be as follows:
The value is positive, again reflecting the increase in disorder going from liquid to vapor. Condensation from vapor to liquid would give a negative value for
Calculations are shown for determining entropy changes at transition temperatures (ice → water or water → vapor and reverse).
Read the material on the link below and answer the following questions:
Is the transfer of heat reversible or irreversible at the transition temperature?
If the phase transition is exothermic, is the entropy change positive or negative?
What is Trouton’s Rule?
What precautions need to be taken in selecting a value for
Why is temperature selection important?
Why would the entropy of vaporization be so much larger than the entropy of fusion?