1. Unit 14: Chemical Thermodynamics
CHM 1046: General Chemistry and Qualitative Analysis
Unit 14: Chemical Thermodynamics
Dr. Jorge L. Alonso
Miami-Dade College – Kendall Campus
Miami, FL
Textbook Reference:
Chapter # 15 (sec.12-17)
Module # 3 (sec. VIII-XII)

2. First Law of Thermodynamics
The law of conservation of energy: energy cannot be created nor destroyed. (James Joule in 1843 )
Energy can, however, be converted from one form to another or transferred from a system to the surroundings or vice versa.
E = q + w
Therefore, the total energy of the universe is a constant.
w = PV
Esys + Esurr = 0
E = q + PV
Esys = -Esurr
E = q + RT

3. Second Law of Thermodynamics
Spontaneity
First Law of Thermodynamics E1
Stone
Do all processes that loose energy occur spontaneously (by themselves, without external influence)??????
+ Work E2
- (work + heat)
Second Law of Thermodynamics
E = E2 – E1

4. Spontaneous Processes
can proceed without any outside intervention.
{Spontaneity}
Processes that are spontaneous in one direction are nonspontaneous in the reverse direction.

5. Spontaneous Processes
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.
T < 0ºC
Is the spontaneity of melting ice dependent on anything?
T > 0ºC

6. Thermodynamics vs. Kinetics
Spontaneity
vs. Speed
C diamond C graphite
Thermodynamics vs. Kinetics

7. Irreversible Processes
Irreversible processes cannot be undone by exactly reversing the change to the system. E1
Stone
Heat energy is lost to dissipation and that energy will not be recoverable if the process is reversed.
+ Work
Spontaneous processes are irreversible.
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. E2
- (work + heat)

8. Entropy (S)
Entropy is a measure of the energy that becomes dissipated and unavailable (friction, molecular motion = heat).
Entropy (S) is a term coined by Rudolph Clausius in the 1850’s. Clausius chose "S" in honor of Sadi Carnot (who gave the first successful theoretical account of heat engines, now known as the Carnot cycle, thereby laying the foundations of the second law of thermodynamics).
Clausius was convinced of the significance of the ratio of heat delivered and the temperature at which it is delivered, q T
Entropy (S) =

9. Entropy (S) ENTROPY Gas Liquid Solid
Entropy can be thought of as a measure of the randomness (disorder) of a system.
It is related to the various modes of motion in molecules.
ENTROPY
Liquid
Like total energy, E, and enthalpy, H, entropy is a state function.
Therefore,
S = Sfinal  Sinitial
Solid
{Entropy.WaterBoiling}

10. Second Law of Thermodynamics
the entropy of the universe increases for spontaneous (irreversible) processes.
Suniv = Ssystem + Ssurroundings > 0
Suniv = Ssystem + Ssurroundings = 0
the entropy of the universe does not change for reversible processes.

11. Second Law of Thermodynamics
All spontaneous processes cause the entropy of the universe to increase.
So what is our fate as a result of the second law operating in our Universe?
ENTROPIC DOOM!

12. Entropy on the Molecular Scale
Molecules exhibit several types of motion (Kinetic energies):
Translational: Movement of a 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.
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 (W) of the thermodynamic system 
Entropy is …….
S = k ln W
…..where k is the Boltzmann constant, 1.38  1023 J/K.

13. Entropy on the Molecular Scale
S = k ln W
…..where k is the Boltzmann constant, 1.38  1023 J/K.
The number of microstates (W) and, therefore, the entropy (S) tends to increase with increases in which variables….
Temperature (T).
Volume (V).
The number of independently moving molecules ().

14. CaCl2 (s) Ca 2+(aq) + 2Cl-(aq)
Entropy Changes
In which of the following does Entropy increase & WHY?…….
Gases are formed from liquids and solids.
Liquids or solutions are formed from solids.
The number of gas molecules (or moles) increases.
Entropy increases with the freedom of motion of molecules.
S(g) > S(l) > S(s)
Heat
H2O (l) H2O(g)
{*Entropy&PhaseOfMatter}
H2O
CaCl2 (s) Ca 2+(aq) + 2Cl-(aq)
{*EntropySolutions.KMnO4(aq)}
Electricity
2 H2O (l) H2 (g) + O2(g)
2 C8H18 (l) + 25 O2 (g)
16 CO2(g) + 18 H2O(g)
 gas= = +9/2 = 4.5 C8H18

15. Third Law of Thermodynamics
The entropy (S) of a pure crystalline substance at absolute zero (-273°C) is 0.

16. Standard Entropies
Standard entropies tend to increase with increasing molar mass.
Larger and more complex molecules have greater entropies (greater ways to execute molecular motions)
{*Entropy&MolecuarSize}
500 K S=921J/nK 200 K

17. Standard Entropies (298 K) from Absolute Entropies (0K)
Absolute Entropy (S)
@ - 237°C (0 K), S = 0
Standard Entropy (S˚)
@ 25°C (298 K), S = ????
Solid
Liquid
Gas
q = mcT
H°vap S
q = mcT S°
H°fus
Calculate the sum of all the infinitesimally small changes in entropy as T varies from T=0  T= 298, by taking its Integral.
q = mcT
298
Temp (K)

18. Entropy Changes in the System
S°syst = T
Entropy changes for a reaction (= system) can be estimated in a manner analogous to that by which H is estimated:
where n and m are the coefficients in the balanced chemical equation.

19. Entropy Changes in the System
Problem: Calculate the standard entropy changes for the following reaction at 25oC.
N2 (g) + 3 H2 (g)  2 NH3 (g)
S° = nS°(prod) - mS°(react)
2(192.5) – [(191.5)+3(130.6)]
S° = J/

20. Thermodynamic Changes in Systems (Chem. Reactions)
Hrxn =  Hf (products) - Hf (reactants) ☺
DGrxn = S DGf (products)  S Gf (reactants)
Appendix 1 (CHM 1046 Module): notice S° is in J not kJ.

21. Entropy Changes in the Surroundings
What in a chemical reaction causes entropy changes in the surroundings?
Heat (q) that flows into or out of the system changes the entropy of the surroundings:
Ssurr ∝ - (qsys)
For an isothermal process: q q
Ssurr=
(qsys) T q q
System
At constant pressure, qsys is simply H for the system. q q
Ssurr=
Hsys T q
Surroundings

22. Entropy Change in the Universe
Problem: Calculate the Suniv for the synthesis of 25oC.
N2 (g) + 3 H2 (g)  2 NH3 (g) H°rxn = kJ/mol
Suniv = Ssyst or rxn Ssurr
nS(prod) - mS(react)
Ssurr =
-Hsys T
2(192.5) – [(191.5)+3(130.6)]
S°syst = J/K·mol
Ssurr = 311 J/K·mol
Suniv = J/K·mol J/K·mol Suniv = J/K·mol

23. Entropy Change in the Universe
Suniv = Ssyst or rxn Ssurr
Ssurr =
-Hsys T
Since:
Then:
Suniv = Ssyst +
Hsystem T
Multiplying both sides by T,
 TSuniv =  TSsyst + Hsyst
TSuniv = Hsyst  TSsyst
TDSuniv is defined as the Gibbs (free) Energy, G.
J. Willard Gibbs USA,

24. Gibbs Free Energy (G) ☺ TSuniv = DGuniv = DHsys  TSsys
When Suniv is positive, G is negative.
When G is negative, the process is spontaneous. ca
Gibbs Energy (-TΔS) measures the "useful" or process-initiating work obtainable from an isothermal, isobaric thermodynamic system. Technically, the Gibbs free energy is the maximum amount of non-expansion work which can be extracted from a closed system or this maximum can be attained only in a completely reversible process.

25. Free Energy Changes ☺ At temperatures other than 25°C,
DG° = DH  TS
How does G change with 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.

26. Spontaneity: Enthalpy & Entropy
DG° = DH  TS
all T
all T
low T
high T

27. Spontaneity: Enthalpy & Entropy
DG = DH( TS)
{Entropy Driven Reactions}
Entropy
H2O
Enthalpy
NH4NO3(s) NH4+(aq) + NO3-(aq)
S = +
H = +
(-TS)
{Enthalpy Driven Reaction}
2 H2(g) + O2 (g) H2O(g)
H = -
S = -
n = 2-3 = -1
(+TS)
{EntropySyst+Surr.FormationOfWater}
{Entropy & Enthalpy Driven Reaction}
H = -
S = +
Na2CO3(s) + HCl(aq) NaCl(aq) + CO2 (g)
(-TS)

28. DG° = DHT(S) Problems DG° = DH  TS 0 = (131.3kJ)  T(.133kJ)
(-763)
– (-804)
+41
(.3549)
– (.2219)
+.1330
Product
Reactant
(-T)
TiCl4(l)  TiCl4(g)
DG° = DH  TS
0 = (131.3kJ)  T(.133kJ)
T = 987

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

30. Standard Free Energy Changes
2 C6H6 (l) + 15 O2 (g)
12 CO2(g) + 6 H2O(g)
Calculate the standard free energy changes for the above 25 °C.
DGrxn = SnDG(prod)  SmG(react) f
Standard Molar Gibbs Energy of Formation (G°f)
CO2 (g)
H2O (g)
C6H6 (l) f
[12 CO2(g) + 6 H2O(g)]
– [2 C6H6 (l) + 15 O2(g)]
[12(-394) + 6(-229)] – [2(125) + 15 (0)]
Grxn = J/mol · K

31. Fourth Law of Thermodynamics: Emergence
Complex emergent systems spontaneously (-G) arise when energy flows through a collection of many interacting particles, resulting in new patterns of complex behaviors that are much more than the sum of the individual parts (+S)
The formation of these complex patterns in emergent systems is more efficient in the dissipation of energy (-H), thus speeds up the increase of entropy in the universe.
G = H -TS
A precise definition of emergence and a useful mathematical formulation of this phenomenon remains elusive.
C  f [n, i, E(t)]
Examples: cells forming living organisms; stars forming galaxies; neurons forming conscious brain.

38. 2002 B

41. 2003 A

43. 2004 A

46. 2004 B

48. 2005 A

50. 2006 (B)

53. 2007 (A)