standard entropy

Examples of standard entropy in the following topics:

• Standard Entropy

  • The standard entropy of a substance (its entropy at 1 atmospheric pressure) helps determine if a reaction will take place spontaneously.
  • The standard entropy of a substance is its entropy at 1 atm pressure.
  • Some typical standard entropy values for gaseous substances include:
  • Scientists conventionally set the energies of formation of elements in their standard states to zero.
  • The standard entropy of reaction helps determine whether the reaction will take place spontaneously.

• Standard Free Energy Changes

  • The standard Gibbs Free Energy is calculated using the free energy of formation of each component of a reaction at standard pressure.
  • These same definitions apply to standard enthalpies and internal energies.
  • Don't confuse these thermodynamic standard states with the "standard temperature and pressure" (STP) widely employed in gas law calculations.
  • To accomplish this, combine the standard enthalpy and the standard entropy of a substance to get the standard free energy of a reaction:
  • The other factor to keep in mind is that enthalpy values are normally given in $\frac{kJ}{mole}$ while entropy values are given in $\frac{J}{K\times mole} $ .

• Solvent Effects

  • From this description of ion formation in solution, it should be clear that both enthalpy and entropy factors will be important to the outcome of an ionization process.
  • Although these two inorganic salts have similar standard enthalpies of solution in water, their standard entropies are quite different.
  • One might expect this entropy change to be positive, since a single molecule in the solid state produces two or more ionic species, accompanied by an increase in system disorder.
  • The overall entropy change for solution of NaCl is positive, reflecting the increased disorder of ionization, but the entropy change for CaF2 solution is strongly negative thanks to the solvation shell structure required by the resulting ions.
  • These different entropy changes are incorporated in the free energy of solution, which is exergonic for NaCl, but endergonic for CaF2.

• Spontaneous and Nonspontaneous Processes

  • The laws of thermodynamics govern the direction of a spontaneous process, ensuring that if a sufficiently large number of individual interactions (like atoms colliding) are involved, then the direction will always be in the direction of increased entropy.
  • The second law of thermodynamics states that for any spontaneous process, the overall ΔS must be greater than or equal to zero; yet, spontaneous chemical reactions can result in a negative change in entropy.
  • The increase in temperature of the reaction surroundings results in a sufficiently large increase in entropy, such that the overall change in entropy is positive.
  • Since the overall ΔS = ΔSsurroundings + ΔSsystem, the overall change in entropy is still positive.
  • An endergonic reaction (also called a nonspontaneous reaction or an unfavorable reaction) is a chemical reaction in which the standard change in free energy is positive, and energy is absorbed.

• Changes in Energy

  • For isolated systems, entropy never decreases.
  • Increases in entropy correspond to irreversible changes in a system.
  • The entropy of a system is defined only if it is in thermodynamic equilibrium.
  • However, the entropy of the system of ice and water has increased more than the entropy of the surrounding room has decreased.
  • Ice melting in a warm room is a common example of increasing entropy.

• Microstates and Entropy

  • With more available microstates, the entropy of a system increases.
  • As a result, entropy (denoted by S) is an expression of disorder or randomness.
  • With more available microstates, the entropy of a system increases.
  • This is the basis of an alternative (and more fundamental) definition of entropy:
  • Therefore, the entropy of a solid is less than the entropy of a liquid, which is much less than the entropy of a gas:

• The Third Law of Thermodynamics and Absolute Energy

  • Specifically, the entropy of a pure crystalline substance at absolute zero temperature is zero.
  • Entropy is related to the number of possible microstates according to $S = k_Bln(\Omega)$, where S is the entropy of the system, kB is Boltzmann's constant, and Ω is the number of microstates (e.g. possible configurations of atoms).
  • The constant value (not necessarily zero) is called the residual entropy of the system.
  • The entropy determined relative to this point (absolute zero) is the absolute entropy.
  • The entropy (S) of a substance (compound or element) as a function of temperature (T).

• Solutions and Entropy Changes

  • Chemists use the term "entropy" to denote this aspect of molecular randomness.
  • Entropy is indeed a fascinating, but somewhat confusing, topic.
  • In a similar manner entropy plays an important role in solution formation.
  • All these factors increase the entropy of the solute.
  • This is the same as saying that the entropy of the solute increases.

• Changes in the Entropy of Surroundings

  • Irreversible reactions result in a change in entropy to the surroundings.
  • The heat from the surroundings (entropy) goes into the ice water and the ice melts.
  • The entropy of the ice water increases while the entropy of the surroundings decreases.
  • Distinguish whether or not entropy of surroundings changes in various reactions

• Vapor Pressure of Nonelectrolyte Solutions

  • When molecules transition from the liquid phase to the gas phase, entropy of the system increases.
  • Entropy of the gaseous state is greater than the entropy of the liquid state because the gaseous molecules occupy a larger volume.
  • If the liquid solvent becomes "diluted" with solute, the entropy of the liquid state increases.
  • Therefore, even though the gaseous state has a higher entropy, the difference in entropy between the two systems decreases.
  • The decrease in entropy difference lowers the vapor pressure.