1. Chapter 18 Entropy, Free Energy and Equilibrium Part II
Dr. Al-Saadi

2. Predicting Spontaneity of a System
18.5
Predicting Spontaneity of a System
Starting from the 2nd law of thermodynamics for a spontaneous process:
ΔSuniv = ΔSsys + ΔSsurr > 0
Also, we have: ΔSsurr = ‒ ΔHsys / T
Substituting ② into ① gives:
ΔSuniv = ΔSsys ‒ > 0
T ΔSuniv = T ΔSsys ‒ ΔHsys > 0
In terms of only the system:
ΔHsys ‒ T ΔSsys < 0 ① ②
ΔHsys T ③
At constant T and P, the process is spontaneous only when the change of entropy and enthalpy of the system is such that ΔHsys ‒ T ΔSsys is less than zero.
Dr. Al-Saadi

3. 18.5
Gibbs Free Energy
What happens when one of the potential driving forces behind a chemical reaction is favorable and the other is not?
In other words, what is the situation when enthalpy and entropy compete with each other?
Gibbs free energy (or simply free energy) is another thermodynamic quantity that reflects the balance between enthalpy and entropy of a system. Gibbs free energy is defined as:
G = H – TS
The change in Gibbs free energy for a system at constant temperature is:
ΔG = ΔH – T ΔS
Dr. Al-Saadi

4. Gibbs Free Energy ΔG = ΔH – T ΔS < 0
18.5
Gibbs Free Energy
Gibbs free energy measures the tendency of a chemical reaction to happen.
ΔG = ΔH – T ΔS
Getting back to equation ③, we have for a spontaneous process:
ΔG = ΔH – T ΔS < 0
ΔG < 0 ; the forward reaction is spontaneous.
ΔG > 0 ; the forward reaction is nonspontaneous.
ΔG = 0 ; the reaction is at equilibrium.
Dr. Al-Saadi

5. 18.5
Gibbs Free Energy
Free energy is another useful thermodynamic function that reflects the balance between ΔH and ΔS.
It is defined as the energy available to do work.
When a particular process is accompanied by a release of usable energy (ΔG is negative), it guarantees that it is spontaneous, without need to consider what happens to the rest of the universe (Like the case we’ve seen for ΔSUniv).
Here are few examples at 25C:
NH3(g) + HCl(g)  NH4Cl(s) ΔG = ‒ 91.1 kJ/mol
spontaneous
½N2(g) + O2(g)  NO2(g) ΔG = kJ/mol
nonspontaneous
Dr. Al-Saadi

6. Predicting the Sign of ΔG
18.5
Predicting the Sign of ΔG
The sign of ΔG can be predicted from knowing the sign of ΔH and ΔS for a reaction.
ΔG = ΔH – T ΔS < 0
Dr. Al-Saadi

7. Calculation Involving ΔG
18.5
Calculation Involving ΔG
Exercise:
The reaction CaCO3(s) CaO(s) + CO2(g) has ΔH = kJ/mol and ΔS = J/K ∙mol. Predict the spontaneity of the reaction at room temperature and at 1000C.
What is the temperature (T) above which the reaction would favor the formation of CaO(s) and CO2(g) ?
ΔG = ΔH – T ΔS
ΔH – T ΔS (or ΔG) < 0
Spontaneous T
ΔH – T ΔS = 0
ΔH – T ΔS (or ΔG) > 0
Nonspontaneous
Dr. Al-Saadi

8. Standard Free-Energy Change
18.5
Standard Free-Energy Change
The standard free energy (ΔGrxn) of a system is the change in free energy when reactants in their standard states are converted to products in their standard states.
Standard states are:
Gases 1 atm [ O2(g) , CO2(g) , CH4(g) ]
Liquids Pure liquid [ H2O (l) , C2H5OH (l) ]
Solids Pure solid [ Na(s) , Mg(s) , AlCl3(s) ]
Elements The most stable form at 1 atm and 25C
Solutions 1 M concentration
Dr. Al-Saadi

9. Standard Free-Energy Change from ΔGof
18.5
Standard Free-Energy Change from ΔGof
Consider the following equation:
a A + b B  c C + d D
The change in standard free-energy (ΔGrxn) is:
ΔGrxn = [c ΔGf (C) + d ΔGf (D)] ‒ [a ΔGf (A) + b ΔGf (B)]
In general, the change of free-energy of a chemical reaction (a system) is given by:
ΔGrxn = Σn ΔGf (products) ‒ Σm ΔGf (reactants)
where n and m are the stoichiometric coefficients of the reactants and products in a given equation.
Dr. Al-Saadi

10. Standard Free Energy of Formation
18.5
Standard Free Energy of Formation
The standard free energy of formation (ΔGf) of a compound is the change in free energy that occurs when 1 mole of that compound is formed from elements in their standard states.
It is very similar to the concept of the standard enthalpy of formation (ΔHf) that you learned in CHEM 101.
N2(g) + 2O2(g) NO2(g) ΔGo = kJ
½ N2(g) + O2(g) NO2(g) ΔGo = 51.8 kJ = ΔGof (NO2)
C(s) + 2H2(g) + ½O2(g) CH3OH(l) ΔGof = –166 kJ/mol
2Fe(s) + ½O2(g) Fe2O3(s) ΔGof = – 741 kJ/mol
standard states
ΔGof = 0
ΔGof = 51.8 kJ/mol
Dr. Al-Saadi

11. Standard Free-Energy Change from ΔGof
18.5
Standard Free-Energy Change from ΔGof
Exercise:
Calculate the standard free-energy change for the following reaction at 25C.
2KClO3(s)  2KCl(s) + 3O2(g)
ΔGrxn = Σn ΔGf (products) ‒ Σm ΔGf (reactants)
= 2[ΔGf (KCl(s)] [ΔGf (O2(g)] – 2[ΔGf (KClO3(s)]
= 2[– kJ/mol] [0] – 2[ – kJ/mol]
= – – (– 579.8) = – kJ/mol
The reaction is spontaneous.
A negative ΔG corresponds to a larger equilibrium constant (K), while a positive ΔG corresponds to a smaller equilibrium constant (K).
Dr. Al-Saadi

12. Standard Free-Energy Change from ΔHof and So
18.5
Standard Free-Energy Change from ΔHof and So
Exercise: Calculate ΔH° and ΔS° for the following dissolution reaction: NH4NO3(s)  NH4+ (aq) + NO3‒ (aq) Use the results of this calculation to determine the value of ΔGo . Does NH4NO3 spontaneously dissolve in water at room temperature?
             ΔHof (kJ/mol)         S°(J/mol ∙ K)
NH4NO3(s)                           
NH4+ (aq)                               113.4
NO3- (aq)                                 146.4
Dr. Al-Saadi

13. Temperature Effect on the Sign of ΔG
18.5
Temperature Effect on the Sign of ΔG
The temperature affects the sign of ΔG of chemical reactions.
Consider the melting of ice:
H2O(s)  H2O(l) ΔH= 6.03 kJ/mol ; ΔS= 22.1 J/K ∙mol
The process is :
spontaneous at 10C. (Ice melts at this temperature)
nonspontaneous at ‒ 10C. (Ice does not melt at this temperature)
at equilibrium at 0C. (Ice and water are at equilibrium at this temperature)
ΔG is ‒ve
ΔG is +ve
ΔG is 0
Dr. Al-Saadi

14. Problems Involving Equilibrium Processes
18.5
Problems Involving Equilibrium Processes
Exercise:
Calculate ΔSsub (in J/K·mol) for the sublimation of iodine in a closed flask at 45°C at equilibrium. ΔHsub = 62.4 kJ/mol.
I2(s)  I2(g)
ΔG = 0 (Equilibrium) ΔH ΔS ΔH T
T =
=>
ΔS =
62.4103 J/mol
( ) K
ΔS =
= 196 J/K·mol
Dr. Al-Saadi

15. Relationship between ΔG and ΔG
18.6
Relationship between ΔG and ΔG
Most of the reactions take place in states other than their standard states. Thus, in order to determine spontaneity of a reaction, we need to know how to calculate ΔG when the reaction is not occurring at standard states.
ΔG = ΔG + RT lnQ
Examples of nonstandard states:
Solutions with concentrations other than 1 M.
Gases having pressures other than 1 atm.
where:
ΔG is the non-standard free energy.
ΔG° is the standard free energy (from appendix 2).
R = J/K ∙ mol
T is in Kelvin.
Q is the reaction quotient. => Q =
[products]
[reactants]
Dr. Al-Saadi

16. Relationship between ΔG and ΔG
18.6
Relationship between ΔG and ΔG
Exercise:
Consider the reaction: H2(g) + Cl2(g)  2HCl(g)
How does the value of ΔGrxn change when the pressures of the H2, Cl2 and HCl gases are changed to 0.25 atm, 0.45 atm and 0.30 atm, respectively, at 25C? (ΔGf for HCl(g) is ‒ kJ/mol)
ΔG = ΔG + RT lnQ
Spontaneous at standard states
Calculate ΔG:
ΔG° = [ 2 ( ‒ kJ/mol)] ‒ [0 + 0] = ‒ kJ/mol
Calculate Q :
Cont.
Dr. Al-Saadi

17. Relationship between ΔG and ΔG
18.6
Relationship between ΔG and ΔG
Find the non-standard free energy.
ΔG = ΔG + RT lnQ
ΔG = ‒ J/mol + (8.314 J/K·mol)(298 K) ln(0.80)
ΔG = ‒ kJ/mol
The reaction becomes even more spontaneous at the given non standard states since ΔG is more negative than ΔG. Thus, more HCl will be formed with the given pressures.
The reaction will continue producing more HCl and consuming more H2 and Cl2 until QP = KP .
[HCl]2
[H2][Cl2]
H2(g) + Cl2(g)  2HCl(g)
Q =
Dr. Al-Saadi

18. Relationship between ΔG and K
18.6
Relationship between ΔG and K
At equilibrium: Q = K and ΔG = 0 Thus, at equilibrium: ΔG = ΔG + RT lnQ 0 = ΔG + RT lnK
The larger the value of K, the more negative ΔG becomes, and the reaction proceeds more towards the product side.
ΔG = ‒ RT lnK
Dr. Al-Saadi

19. Relationship between ΔG and K
18.6
Relationship between ΔG and K
ΔG = ‒ RT lnK
At equilibrium, there is a significant conversion of reactants to products.
Dr. Al-Saadi

20. Relationship between ΔG and K
18.6
Relationship between ΔG and K
ΔG = ‒ RT lnK
At equilibrium, there is a significant conversion of products to reactants.
Dr. Al-Saadi

21. Relationship between ΔG and K
18.6
Relationship between ΔG and K
Example:
Using the table of standard free energies, calculate the equilibrium constant, KP, for the following reaction at 25C.
2HCl(g) H2(g) Cl2(g)
ΔG = [0 + 0] ‒ [ 2 ( ‒ kJ/mol)]
= kJ/mol (nonspontaneous)
Substitute into:
ΔG = ‒ RT lnKP
× 103 J/mol = ‒ (8.314 J/K·mol)(298 K) ln KP
KP = 3.98 x 1034
K << 1 (reactants are favored)
Dr. Al-Saadi

22. Relationship between ΔG and K
18.6
Relationship between ΔG and K
Example:
The overall reaction for corrosion is:
4Fe (s) + 3O2(g) Fe2O3(s)
The ΔHf for Fe2O3 (s) is ‒ 826 kJ/mol, and S (J/K·mol) for Fe2O3(s), Fe(s) and O2(g) are 90, 27 and 205, respectively. Calculate K for this reaction at 25C.
Dr. Al-Saadi

23. Thermodynamics of Living Systems
18.7
Thermodynamics of Living Systems
Thermodynamics have a great effect in biological sciences, such as processes taking place inside our bodies.
Many chemical reactions carried out inside the body (such as DNA and protein formation) are not spontaneous, but they can proceed through coupled reactions.
Proteins are polymers made from connecting different amino acids together. However, formation of proteins from amino acids is a nonspontaneous process.
For example:
alanine + glycine alanylglycine
ΔG = + 29 kJ/mol
Amino acids
Protein molecule
Dr. Al-Saadi

24. Thermodynamics of Living Systems
18.7
Thermodynamics of Living Systems
Tracking chemistry inside our bodies:
Metabolism process:
C6H12O6 (s) + 6O2 (g)  6CO2 (g) + 6H2O(l) ΔG = ‒ 2880 kJ/mol
The free energy released from metabolism is used to synthesize energy-storage molecules called adenosine triphosphate (ATP):
Based on the body needs, ATP are hydrolyzed back to ADP and releases 31 kJ/mol of free energy, that can be utilized to drive unfavorable (nonspontaneous) important chemical reactions, such as protein synthesis.
+ H2O
ΔG = + 31 kJ/mol
+ H3PO4
ADP
ATP
Dr. Al-Saadi

25. Thermodynamics of Living Systems
18.7
Thermodynamics of Living Systems
Tracking chemistry inside our bodies:
With the help of enzymes, the nonspontaneous protein synthesis reaction inside the body is coupled to the spontaneous ATP hydrolysis to favor the formation of proteins as follows:
alanine + glycine  alanylglycine ΔG = + 29 kJ/mol
ATP H2O (l)  ADP H3PO4 (aq) ΔG = ‒ 31 kJ/mol
alanine + glycine + ATP + H2O (l)  ADP + H3PO4 (aq) + alanylglycine
ΔG = ‒ 2 kJ/mol
Dr. Al-Saadi

26. Exercises For which process is S negative?
1) Evaporation of 1 mol of CCl4(l).
2) Mixing 5 mL ethanol with 25 mL water.
3) Compressing 1 mol Ne at constant temperature from 1.5 atm to 0.5 atm.
4) Raising the temperature of 100 g Cu from 275 K to 295 K.
5) Grinding a large crystal of KCl to powder. 
Dr. Al-Saadi

27. Exercises Which of the following shows a decrease in entropy?
1) Precipitation
2) Gaseous reactants forming a liquid
3) A burning piece of wood
4) Melting ice
5) Two of these   
Dr. Al-Saadi

28. Exercises
At 1 atm, liquid water is heated above 100°C. Ssurr for this process is:
1) greater than zero.
2) less than zero.
3) equal to zero.
4) more information is needed to answer this question. 
Dr. Al-Saadi

29. Exercises Given the following free energies of formation: Gf
C2H2(g) kJ/mol
C2H6(g) –32.9 kJ/mol
 Calculate Kp at 298 K for:
C2H2(g) + 2H2(g) C2H6(g)
1) 9.07  10–1
2) 97.2
3) 1.24  1031
4) 2.72  1042
5) 7.55  1051 
Dr. Al-Saadi