1. Chapter 19 Chemical Thermodynamics

2. 19.1 Spontaneous Processes
The first law of thermodynamics states that energy is conserved, i.e., it cannot be created or destroyed in any process.
Therefore, the total energy of the universe is a constant.
Energy can, however, be converted from one form to another or transferred from a system to the surroundings or vice versa.
Mathematically as Δ E = q + w for a process or reaction.
The first law doesn’t address the extent or direction of a process or reaction.

3. Spontaneous processes are those that proceed on their own without any outside assistance.
A spontaneous process has a definite direction.
E.g., the gas in vessel B will spontaneously effuse into vessel A, but once the gas is in both vessels it will not spontaneously return to its original state.

4. Processes that are spontaneous in one direction are nonspontaneous in the reverse direction.

5. Processes that are spontaneous at one temperature may be nonspontaneous at other temperatures.
Above 0C it is spontaneous for ice to melt.
Below 0C the reverse process is spontaneous.

6. A Criterion for Spontaneity
Spontaneous reactions can be both exothermic and endothermic in terms of energy flow.
So, what is the factor at work in determining the natural direction of these processes?
Understanding spontaneity lies in distinguishing between reversible and irreversible paths between states of a system

7. Reversible and Irreversible Processes
Entropy (S) is a term coined by Rudolph Clausius in 1824.
Clausius was convinced of the significance of the ratio of heat delivered and the temperature at which it is delivered, q T

8. Reversible Processes:
In a reversible process the system changes in such a way that the system and surroundings can be put back in their original states by exactly reversing the process.

9. Irreversible Processes:
Irreversible processes cannot be undone by exactly reversing the change to the system.
Spontaneous processes are irreversible.

10. Entropy can be thought of as a measure of the randomness of a system.
It is related to the various modes of motion in molecules.
Like total energy, E, and enthalpy, H, entropy is a state function.
Therefore,
S = Sfinal  Sinitial

11. Entropy Change qrev T S =
For a process occurring at constant temperature (an isothermal process), the change in entropy is equal to the heat that would be transferred if the process were reversible divided by the temperature:
S =
qrev T

12. The Second Law of Thermodynamics
The second law of thermodynamics states that the entropy of the universe increases for spontaneous processes, and the entropy of the universe does not change for reversible processes.
In other words:
For reversible processes:
Suniv = Ssystem + Ssurroundings = 0
For irreversible processes:
Suniv = Ssystem + Ssurroundings > 0

13. These last truths mean that as a result of all spontaneous processes the entropy of the universe increases.

14. 19.3 The Molecular Interpretation of Entropy
Ludwig Boltzmann, a 19th c Austrian physicist, described the concept of entropy on the molecular level.
Temperature is a measure of the average kinetic energy of the molecules in a sample.

15. Molecules exhibit several types of motion:
Translational: Movement of the entire molecule from one place to another.
Vibrational: Periodic motion of atoms within a molecule.
Rotational: Rotation of the molecule on about an axis or rotation about  bonds. 

16. Boltzmann envisioned the motions of a sample of molecules at a particular instant in time.
This would be akin to taking a snapshot of all the molecules.
He referred to this sampling as a microstate of the thermodynamic system.

17. Each thermodynamic state has a specific number of microstates, W, associated with it.
Entropy is
S = k lnW
where k is the Boltzmann constant, 1.38  1023 J/K.

18. The change in entropy for a process, then, is
S = k lnWfinal  k lnWinitial
lnWfinal
lnWinitial
S = k ln
Entropy increases with the number of microstates in the system.

19. The number of microstates and, therefore, the entropy tends to increase with increases in
Temperature.
Volume.
The number of independently moving molecules.

20. Making Qualitative Predictions about ΔS
Entropy increases with the freedom of motion of molecules.
Therefore,
S(g) > S(l) > S(s)
An increase in the number of microstates and, hence entropy, parallels an increase in
Temperature
Volume
Number of independently moving particles
In ice, molecular motion is restricted to vibrations.
As water melts, molecules also move about with translational and rotation motion, increasing the number microstates and entropy.

21. Generally, when a solid is dissolved in a solvent, to form a solution, entropy increases because the ions have more motional energy dissolved than when solid.
However, water molecules attracted by ion-dipole attractions decreases motional for water molecules.
Sometimes the dissolving of salts that contain highly-charged ions can result in a net decrease in entropy.

22. In general, entropy increases when
The fewer independently moving particles resulting from some reactions (e.g., as shown) results in fewer degrees of freedom and a decrease in entropy.
In general, entropy increases when
Gases are formed from liquids and solids.
Liquids or solutions are formed from solids.
The number of gas molecules increases.
The number of moles increases.

23. Third Law of Thermodynamics
The entropy of a pure crystalline substance at absolute zero is 0.
Vibrational motion caused when temperature is increased causes increase in possible microstates and an increase in entropy.
See Fig showing entropy as a function of temperature.

24. 19.4 Entropy Changes in Chemical Reactions
Entropy changes for a reaction can be estimated in a manner analogous to that by which H is estimated:
S° = nS°(products) - mS°(reactants)
where n and m are the coefficients in the balanced chemical equation.
The theory and methods for these calculations are beyond the scope of your text.

25. These are standard molar entropy (So) values of substances in their standard states.

26. Standard entropies tend to increase with increasing molar mass.
Larger and more complex molecules have greater entropies.

27. Observations based on standard molar entropy (So) values.
Unlike enthalpies of formation, standard molar entropies for elements are not zero.
Standard molar entropies of gases are greater than those for liquids and solids (Fig 19.14).
Standard molar entropies generally increase with increasing molar mass.
Standard molar entropies generally increase with an increasing number of atoms in the formula of a substance.

28. Sample problem:

29. Entropy Changes in Surroundings
Heat that flows into or out of the system changes the entropy of the surroundings.
For an isothermal process:
Ssurr =
qsys T
At constant pressure, qsys is simply H for the system.

30. Entropy Change in the Universe
The universe is composed of the system and the surroundings.
Therefore,
Suniverse = Ssystem + Ssurroundings
For spontaneous processes
Suniverse > 0

31. Entropy Change in the Universe
This becomes:
Suniverse = Ssystem +
Multiplying both sides by T,
TSuniverse = Hsystem  TSsystem
Hsystem T

32. 19.5 Gibbs Free Energy
A thermodynamic state function that combines enthalpy and entropy in the form G = H – TS
For a change occurring at constant temperature and pressure, the change in free energy is
ΔG = ΔH - TΔS
TDSuniverse is defined as the Gibbs free energy, G.
When Suniverse is positive, G is negative.
Therefore, when G is negative, a process is spontaneous.

33. If DG is negative, the forward reaction is spontaneous.
If DG is 0, the system is at equilibrium.
If G is positive, the reaction is spontaneous in the reverse direction.

34. Standard Free Energy Changes
Analogous to standard enthalpies of formation are standard free energies of formation, G. f
DG = SnDG(products)  SmG(reactants) f
where n and m are the stoichiometric coefficients.

35. At temperatures other than 25°C,
DG° = DH  TS
How does G change with temperature?

36. 19.6 Free Energy and Temperature
There are two parts to the free energy equation:
H— the enthalpy term
TS — the entropy term
The temperature dependence of free energy, then comes from the entropy term.

38. 19.7 Free Energy and Equilibrium
Under any conditions, standard or nonstandard, the free energy change can be found this way:
G = G + RT lnQ
(Under standard conditions, all concentrations are 1 M, so Q = 1 and lnQ = 0; the last term drops out.)

39. At equilibrium, Q = K, and G = 0.
The equation becomes
0 = G + RT lnK
Rearranging, this becomes
G = RT lnK
or,
K = eG/RT