1. Thermodynamics
Part 5 - Spontaneity

2. ΔHrxn = nΔHproducts - nΔHreactants
Thermodynamics
Thermodynamics = the study of energy changes that accompany physical and chemical changes.
Enthalpy (H): the total energy “stored” within a substance
Enthalpy Change (ΔH): a comparison of the total enthalpies of the product & reactants.
ΔHrxn = nΔHproducts - nΔHreactants

3. Exothermic vs. Endothermic
Exothermic reactions/changes: release energy in the form of heat; have negative ΔH values.
H2O(g)  H2O(l) ΔH = kJ
Endothermic reactions/changes: absorb energy in the form of heat; have positive ΔH values.
H2O(l)  H2O(g) ΔH = kJ

4. Reaction Pathways Endothermic Exothermic
Changes that involve a decrease in enthalpy are favored!
Endothermic
Exothermic Ea Ea
energy
energy P R R P
time
time
Ea = activation energy; P = products; R = reactants

5. ΔSrxn = nΔ Sproducts – nΔSreactants
Entropy
Entropy (S): the measure of the degree of disorder in a system; in nature, things tend to increase in entropy, or disorder.
ΔSrxn = nΔ Sproducts – nΔSreactants
All physical & chemical changes involve a change in entropy, or ΔS. (Remember that a high entropy is favorable)

6. Entropy

7. Entropy

8. Entropy
gases
solids
liquids
solutions

9. Entropy
pure substances
mixtures

10. Entropy

11. Entropy

12. Driving Forces in Reactions
Enthalpy and entropy are DRIVING FORCES for spontaneous reactions (rxns that happen at normal conditions)
It is the interplay of these 2 driving forces that determines whether or not a physical or chemical change will actually happen.

13. Free Energy
Free Energy (G): relates enthalpy and entropy in a way that indicates which predominates; the quantity of energy that is available or stored to do work or cause change.

14. Free Energy ΔG = ΔH – TΔS Where: ΔG = change in free energy (kJ)
ΔH = change in enthalpy (kJ)
T = absolute temp (K)
ΔS = change in entropy (kJ/K)

15. Free Energy ΔG: positive (+) value means change is NOT spontaneous
ΔG: negative (-) value means change IS spontaneous

16. Relating Enthalpy and Entropy to Spontaneity
Example ΔH ΔS
Spontaneity
2K + 2H2O  2KOH + H2 - +
always spon.
H2O(g)  H2O(l)
lower temp.
H2O(s)  H2O(l)
higher temp.
16CO2+18H2O2C8H18+25O2
never spon.

17. Example #1 For the decomposition of O3(g) to O2(g): 2O3(g)  3O2(g)
ΔH = kJ
ΔS = °C
a) Calculate ΔG for the reaction.
ΔG = ( kJ) – (298K)( kJ/K)
ΔG = kJ

18. Example #1 YES For the decomposition of O3(g) to O2(g):
2O3(g)  3O2(g)
ΔH = kJ
ΔS = °C
b) Is the reaction spontaneous?
YES

19. Example #1 For the decomposition of O3(g) to O2(g): 2O3(g)  3O2(g)
ΔH = kJ
ΔS = °C
c) Is ΔH or ΔS (or both) favorable for the reaction?
Both ΔS and ΔH are favorable
(both are driving forces)

20. Example #2 Fe2O3(s) + 3CO(g)  2Fe(s) + 3CO2(g) ΔG = ΔH – TΔS
What is the minimum temperature (in °C) necessary for the following reaction to occur spontaneously?
Fe2O3(s) + 3CO(g)  2Fe(s) + 3CO2(g)
ΔH = kJ; ΔS = J/K
(Hint: assume ΔG = kJ)
ΔG = ΔH – TΔS
= (144.5) – (T)(0.0243)
T ≈ 5950 K
T = 5677 °C