1. Chapter 17 Free Energy and Thermodynamics

2. Thermodynamics vs. Kinetics

3. CHEMICAL THERMODYNAMICS The first law of thermodynamics:
Energy and matter can be neither created nor destroyed; only transformed from one form to another. The energy and matter of the universe is constant.
The second law of thermodynamics:
In any spontaneous process there is always an increase in the entropy of the universe.
The entropy is increasing.
The third law of thermodynamics:
The entropy of a perfect crystal at 0 K is zero.
There is no molecular motion at absolute 0 K.

4. H = Enthalpy = heat of reaction = qp
STATE FUNCTIONS
A property of a system which depends only on its present state and not on its pathway.
H = Enthalpy = heat of reaction = qp
A measure of heat (energy) flow of a system relative to its surroundings.
H° standard enthalpy
Hf° enthalpy of formation
H° = n Hf° (products) -  m Hf° (reactants)
H = U + PV
U represents the Internal energy of the particles, both the kinetic and potential energy. U = q + w

5. HEAT VS WORK energy transfer as a energy expanded to
result of a temperature move an object against
difference a force
qp w = F x d
endothermic (+q) work on a system
(+w)
exothermic (-q) work by the system
(-w)
qc = -qh w = -PV

6. SPONTANEOUS PROCESSES A measure of randomness or disorder in a system.
A spontaneous process occurs without outside intervention. The rate may be fast or slow.
Entropy
A measure of randomness or disorder in a system.
Entropy is a state function with units of J/K and it can be
created during a spontaneous process.
Suniv = Ssys + Ssurr
The relationship between Ssys and Ssurr
Ssys Ssurr Suniv Process spontaneous?
Yes
No (Rx will occur in
opposite direction)
? Yes, if Ssys > Ssurr
? Yes, if Ssurr > Ssys

7. Comparing Potential Energy
The direction of spontaneity can be determined by comparing the potential energy of the system at the start and the end.

8. Reversibility of Process
any spontaneous process is irreversible
it will proceed in only one direction
a reversible process will proceed back and forth between the two end conditions
equilibrium
results in no change in free energy
if a process is spontaneous in one direction, it must be nonspontaneous in the opposite direction

9. Diamond → Graphite
Graphite is more stable than diamond, so the conversion of diamond into graphite is spontaneous – but don’t worry, it’s so slow that your ring won’t turn into pencil lead in your lifetime (or through many of your generations).

10. Predicting Relative S0 Values of a System
1. Temperature changes
S0 increases as the temperature rises.
2. Physical states and phase changes
S0 increases as a more ordered phase changes to a less ordered phase.
3. Dissolution of a solid or liquid
S0 of a dissolved solid or dissolved liquid is usually greater than the S0 of the pure solute. However, the extent depends upon the nature of the solute and solvent.
4. Dissolution of a gas
A gas becomes more ordered when it dissolves in a liquid or solid.
5. Atomic size or molecular complexity
In similar substances, increases in mass relate directly to entropy.
In allotropic substances, increases in complexity (e.g. bond flexibility) relate directly to entropy.

11. Factors Affecting Whether a Reaction Is Spontaneous
The two factors that determine the thermodynamic favorability are the enthalpy and the entropy.
The enthalpy is a comparison of the bond energy of the reactants to the products.
bond energy = amount needed to break a bond. DH
The entropy factors relates to the randomness/orderliness of a system DS
The enthalpy factor is generally more important than the entropy factor

12. S0rxn = S0products - S0reactants
Entropy Changes in the System
S0rxn - the entropy change that occurs when all reactants
and products are in their standard states.
S0rxn = S0products - S0reactants
The change in entropy of the surroundings is directly
related to an opposite change in the heat of the system
and inversely related to the temperature at which the
heat is transferred.
Ssurroundings = -
Hsystem T

13. Increases in Entropy

14. Entropy S = Sf - Si S > q/T S = H/T
For a reversible (at equilibrium) process
H - T  S < 0
For a spontaneous reaction at constant T & P
 H - T S
If the value for  H - T S is negative for a reaction then the reaction is spontaneous in the direction of the products.
If the value for  H - T S is positive for a reaction then the reaction is spontaneous in the direction of the reactants. (nonspontaneous for products)

15. A spontaneous endothermic chemical reaction.
water
Ba(OH)2 8H2O(s) + 2NH4NO3(s) Ba2+(aq) + 2NO3-(aq) + 2NH3(aq) + 10H2O(l) .
DH0rxn = kJ

16. Predicting the sign of S° The sign is positive if:
APPLICATION OF THE 3RD LAW OF THERMODYNAMICS
S° = standard entropy = absolute entropy
S is usually positive (+) for Substances, S can be negative (-) for Ions because H3O+ is used as zero
Predicting the sign of S°
The sign is positive if:
1. Molecules are broken during the Rx
2. The number of moles of gas increases
3. solid  liquid liquid  gas solid  gas
an increase in order occurs
1. Ba(OH)2•8H2O (s) + 2NH4NO3(s) 2NH3(g) + 10H2O(l) + Ba(NO3)2(aq)
2. 2SO2(g) + O2(g)  2SO3(g)
3. HCl(g) + NH3(g)  NH4Cl(s)
4. CaCO3(s)  CaO(s) + CO2(g)

17. Ex. 17.2a – The reaction C3H8(g) + 5 O2(g)  3 CO2(g) + 4 H2O(g) has DHrxn = kJ at 25°C. Calculate the entropy change of the surroundings.
Given:
Find:
DHsystem = kJ, T = 298 K
DSsurroundings, J/K
Concept Plan:
Relationships: DS
T, DH
Solution:
Check:
combustion is largely exothermic, so the entropy of the surrounding should increase significantly

18.  S°=  n S°(product)- m S°(reactant)
1. Acetone, CH3COCH3, is a volatile liquid solvent. The standard enthalpy of formation of the liquid at 25°C is kJ/mol; the same quantity for the vapor is kJ/mol. What is Ssurr when 1.00 mol liquid acetone vaporizes at 25oC?
Calculate S° at 25° for:
2 NaHCO3 (s)  CO2 (g) + Na2CO3 (s) + H2O (g)
Table of Sf

19. DG0rxn = S mDG0products - S nDG0reactants
Gibbs Free Energy (G)
G, the change in the free energy of a system, is a measure of the spontaneity of the process and of the useful energy available from it. It is also a State Function just like entropy and enthalpy.
DG0rxn = S mDG0products - S nDG0reactants

20. STANDARD FREE ENERGY OF FORMATION
G°f
The free energy change that occurs when 1 mol of substance is formed from the elements in their standard state.
Calculate G° at 25oC for:
2 CH3OH(l) + 3 O2(g)  2 CO2(g) + 4 H2O(g)
Table of DGf

21. Workshop T1 on Entropy & Gibbs Free Energy
Carbon tetrachloride is a volatile liquid. What is Ssurr when 1.00 mol liquid carbon tetrachloride vaporizes at 25oC?
Calculate S° and G° at 25° for:
a. CaCO3 (s)  CO2 (g) + CaO (s)
2 AgBr(s) + I2(s)  Br2(l) AgI(s)
Table of Sf

22. DG0system = DH0system - TDS0system
Gibbs Free Energy (G)
G, the change in the free energy of a system, is a measure of the spontaneity of the process and of the useful energy available from it. There are a multiple number of ways to calculate DG.
DG0system = DH0system - TDS0system
G < 0 for a spontaneous process
G > 0 for a nonspontaneous process
G = 0 for a process at equilibrium

23. Free Energy Change and Spontaneity

24. INTERPRETING G° FOR SPONTANEITY
1. When G° is very small (less than -10 KJ) the reaction is spontaneous as written. Products dominate.
G° < 0 G°(R) > G°(P)
2. When G° is very large (greater than 10 KJ) the reaction is non spontaneous as written. Reactants dominate.
G° > 0 G°(R) < G°(P)
3. When G° is small (+ or -) at equilibrium then both
reactants and products are present.
G° = 0
Ba(OH)•8 H2O(s) + 2 NH4NO3(s)  2 NH3(g)+10 H2O(l) + Ba(NO3)3(aq)

25. G = H - TS describes the temperature dependence of spontaneity
GIBBS FREE ENERGY : G
G = H - TS
describes the temperature dependence of spontaneity
Standard conditions (1 atm, if soln=1M & 25°):
G° = H° - TS°
A process (at constant P & T) is spontaneous in the direction in which the free energy decreases.

26. Ex. 17.3a – The reaction CCl4(g)  C(s, graphite) + 2 Cl2(g) has DH = kJ and DS = J/K at 25°C. Calculate DG and determine if it is spontaneous.
Given:
Find:
DH = kJ, DS = J/K, T = 298 K
DG, kJ
Concept Plan:
Relationships: DG
T, DH, DS
Solution:
Since DG is +, the reaction is not spontaneous at this temperature. To make it spontaneous, we need to increase the temperature.
Answer:

27. Ex. 17.3a – The reaction CCl4(g)  C(s, graphite) + 2 Cl2(g) has DH = kJ and DS = J/K. Calculate the minimum temperature it will be spontaneous.
Given:
Find:
DH = kJ, DS = J/K, DG < 0
T, K
Concept Plan:
Relationships: T
DG, DH, DS
Solution:
The temperature must be higher than 673K for the reaction to be spontaneous
Answer:

28. Workshop T2 2 SO2(g) + O2(g)  2 SO3(g)
1. At 25°C & 1 atm, calculate H°, S° & G° using the tables found in the appendix.
2. Calculate DG at 25°C & 1 atm using G°=H°-TS°, compare the value obtained in question 1.
3. Calculate the minimum temperature that the above reaction becomes spontaneous.

29. G AND EQUILIBRIUM
The equilibrium point occurs at the lowest free energy available to the reaction system.
When a substance undergoes a chemical reaction, the reaction proceeds to give the minimum free energy at equilibrium.
G = G° + RT ln (Q)
but at equilibrium: G = 0
G° = 0 then K = 1
G° < 0 then K > 1
G° > 0 then K < 1
Q: Corrosion of iron by oxygen is
4 Fe(s) O2(g)  2 Fe2O3(s)
calculate K for this Rx at 25°C.
G = G° + RT ln (Q) nonstandard conditions
G° = -RT ln (K) standard conditions

31. Table 1 The Relationship Between DG0 and K at 250C
DG0(kJ) K
Significance
Essentially no forward reaction; reverse reaction goes to completion
REVERSE REACTION
200
9x10-36
FORWARD REACTION
100
3x10-18 50
2x10-9 10
2x10-2 1
7x10-1
Forward and reverse reactions proceed to same extent 1 -1
1.5
-10
5x101
-50
6x108
Forward reaction goes to completion; essentially no reverse reaction
-100
3x1017
-200
1x1035

32. DG = +46400 J + (8.314 J/K)(700 K)(ln 1.2 x 10-7)
Example - DG
Calculate DG at 427°C for the reaction below if the PN2 = 33.0 atm, PH2= 99.0 atm, and PNH3= 2.0 atm
N2(g) + 3 H2(g) ® 2 NH3(g)
Q =
PNH32
PN21 x PH23
(2.0 atm)2
(33.0 atm)1 (99.0)3 =
= 1.2 x 10-7
DH° = [ 2(-46.19)] - [0 +3( 0)] = kJ = J
DS° = [2 (192.5)] - [(191.50) + 3(130.58)] = J/K
DG° = J - (700 K)( J/K)
DG° = J
DG = DG° + RTlnQ
DG = J + (8.314 J/K)(700 K)(ln 1.2 x 10-7)
DG = J = -46 kJ

33. Free Energy, Equilibrium and Reaction Direction
If Q/K < 1, then ln Q/K < 0; the reaction proceeds to the right (DG < 0)
If Q/K > 1, then ln Q/K > 0; the reaction proceeds to the left (DG > 0)
If Q/K = 1, then ln Q/K = 0; the reaction is at equilibrium (DG = 0)
DG = RT ln Q/K = RT lnQ - RT lnK
Under standard conditions (1M concentrations, 1atm for gases), Q = 1 and ln Q = 0 so
DG0 = - RT lnK

34. Example - K
Estimate the equilibrium constant and position of equilibrium for the following reaction at 427°C
N2(g) + 3 H2(g) Û 2 NH3(g)
DH° = [ 2(-46.19)] - [0 +3( 0)] = kJ = J
DS° = [2 (192.5)] - [(191.50) + 3(130.58)] = J/K
DG° = J - (700 K)( J/K)
DG° = J
*estimation*
DG° = -RT lnK
J = -(8.314 J/K)(700 K) lnK
lnK = -7.97
K = e-7.97 = 3.45 x 10-4
since K is << 1, the position of equilibrium favors reactants

35. AgI (s)  Ag+ (aq) + I- (aq) BaSO4 (s) Ba2+(aq) + SO42-(aq)
Workshop T3 on DGo & K
Calculate Gº at 25ºC
AgI (s)  Ag+ (aq) + I- (aq)
BaSO4 (s) Ba2+(aq) + SO42-(aq)
What is the value for Ksp at 25ºC?
Calculate K at 25ºC for
Zn(s) + 2H+(aq) Zn2+(aq) + H2 (g).
Table

36. Temperature Dependence of K
for an exothermic reaction, increasing the temperature decreases the value of the equilibrium constant
for an endothermic reaction, increasing the temperature increases the value of the equilibrium constant

37. Gº & Spontaneity is dependent on Temperature
Hº Sº Gº
Spontaneous at all T
Non spontaneous at all T
/- At Low T= Spontaneous
At High T= Nonspontaneous
/- At low T= Nonspontaneous
At High T= Spontaneous

38. Lecture Problems on Temperature Relationships
1. At what temperature is the following process spontaneous at 1 atm?
Br2 (l)  Br2 (g)
What is the normal boiling point for Br2 (l)?
2. Calculate Go & Kp at 35ºC
N2O4 (g)  2 NO2 (g)

39. Workshop T4 on Temperature Relationships
1. Predict the Spontaneity for H2O(s)  H2O(l)
at (a) -10ºC , (b) 0ºC & (c) 10ºC.
2. Carbon tetrachloride is a volatile liquid. What is its normal boiling point?
3. Calculate Go & Kp at 375ºC
4CO (g) + 8H2 (g)  4CH4OH (g)
4. Calculate Hº, Sº & Gº at 25ºC and 650ºC.
CS2 (g) + 4H2 (g) CH4 (g) + 2H2S(g)
Compare the two values and briefly discuss the spontaneity of the Rx at both temperature.