Standard molar entropy
In chemistry, the standard molar entropy is the entropy content of one mole of pure substance at a standard state of pressure and any temperature of interest. These are often (but not necessarily) chosen to be the standard temperature and pressure.
The standard molar entropy at pressure = is usually given the symbol S°, and has units of
Thermodynamics
If a mole of a solid substance is a perfectly ordered solid at 0 K, then if the solid is warmed by its surroundings to 298.15 K without melting, its absolute molar entropy would be the sum of a series of N stepwise and reversible entropy changes. The limit of this sum as becomes an integral:
In this example, and is the molar heat capacity at a constant pressure of the substance in the reversible process k. The molar heat capacity is not constant during the experiment because it changes depending on the (increasing) temperature of the substance. Therefore, a table of values for is required to find the total molar entropy. The quantity represents the ratio of a very small exchange of heat energy to the temperature T. The total molar entropy is the sum of many small changes in molar entropy, where each small change can be considered a reversible process.
Chemistry
The standard molar entropy of a gas at STP includes contributions from:[2]
- The heat capacity of one mole of the solid from 0 K to the melting point (including heat absorbed in any changes between different crystal structures).
- The latent heat of fusionof the solid.
- The heat capacity of the liquid from the melting point to the boiling point.
- The latent heat of vaporizationof the liquid.
- The heat capacity of the gas from the boiling point to room temperature.
Changes in entropy are associated with
The standard entropy of reaction helps determine whether the reaction will take place spontaneously. According to the second law of thermodynamics, a spontaneous reaction always results in an increase in total entropy of the system and its surroundings:
Molar entropy is not the same for all gases. Under identical conditions, it is greater for a heavier gas.
See also
References
- ^ Pauling, Linus (1960). The Nature of the Chemical Bond (3rd ed.). Ithaca, NY: Cornell University Press.
- ISBN 1-889526-15-0.
- ISBN 0-07-251264-4.