1. ENTHALPY, ENTROPY AND GIBBS FREE ENERGY

2. The First Law of Thermodynamics
Energy can neither be created or destroyed
The energy of the universe is constant, but it can change forms.

3. Energy book keeper
First Law accounts for energy, but it does not tell us why a particular process occurs in a given direction

4. Spontaneity DOES NOT MEAN FAST!!!
Means that the process occurs without any outside intervention

5. Energy released or absorbed during a chemical reaction (heat of reaction) is equal to the difference between the potential energy of the products and the potential energy of the reactants.
In a chemical reaction; reactants --> products
ΔPE = PE products - PE reactants
PE can be thought as heat energy (H)
Therefore, ΔH (kJ) = H products - H reactants

6. When ΔH is negative H products < H reactants and the reaction is exothermic.

7. When ΔH is positive H products > H reactants and the reaction is endothermic.

8. Energy released or absorbed by a chemical reaction can be represented by a potential energy diagram.

9. Activation Energy
The activation energy is the minimum energy required to start a chemical reaction by providing colliding molecules with enough energy for effective collisions to occur.
The activated complex is the short-lived and unstable intermediate species located at the highest of the activation energy.    

10. Catalysts
A catalyst provides an alternate reaction pathway, which has a lower activation energy than an uncatalyzed reaction.

11. Look at Table A-6 Notice the following:
Substances are in alphabetical order
ΔHf (enthalpy of formation) is in kJ/mole
Free elements have a ΔHf = 0
(they are not compounds formed from elements)
The enthalpy of reaction is equal to the sum of the enthalpies of formation for the products – the sum of the enthalpies of formation for the reactants Σ ΔHr = ΔHf products - Σ ΔHf reactants

12. ENTHALPY CALCULATIONS
2CO(g) O2(g)  2CO2(g)
ΔHr = Σ ΔHf products - Σ ΔHf reactants
=[2 mole( kJ/mole)] -
[2mole( kJ/mole) + 1mole(0kJ/mole)]
= [ kJ] - [ kJ]
= kJ
The reaction is exothermic

13. ENTROPY Entropy is the degree of disorder
Represented with the symbol S
Matter changes from a more ordered to less ordered state
2H2O  2H O2
H2(l)  H2(g)
A positive ∆S means an increase in entropy

14. States of matter
Ssolid < Sliquid << Sgas

15. Which has more entropy? 1. Solid or gaseous phosphorus
2. CH4(g) or C3H8(g)
3. NaCl(s) or NaCl(aq)

16. Second Law of Thermodynamics
In any spontaneous process there is always an increase in the entropy of the universe
The entropy of the universe is constantly increasing

17. ENTROPY CALCULATIONS
If the reaction increases entropy, ∆S is positive and the reaction is said to be ENTROPY-FAVORED
Calculate the entropy change(∆S) for the following reaction
CH4(g) O2(g)  CO2(g) H2O(l)

18. Three S’s
Ssys = system
Ssurr = surroundings
Ssys + Ssurr = Suniv

19. Suniv If it is +, the entropy of the universe is increasing
Process is spontaneous
If it is negative, the process is not spontaneous

20. Change of state H2O(l)  H2O(g) What happens to the S of the water?
Ssys= +

21. What about surroundings?
Heat is flowing from the surroundings to the system
Random motion of particles decreases
Ssurr = -

22. Which S controls the situation? DEPENDS ON TEMP
Is it spontaneous?
Need to look at Suniv
Which S controls the situation?
DEPENDS ON TEMP

23. Exothermic Process Always increases entropy of surroundings
But, its significance depends on the temp at which the process occurs
Energy transfer will be more significant at lower temps

24. GIBBS FREE ENERGY CH4(g) + 2O2(g)  CO2(g) + 2H2O(l)
Gibbs free energy (∆G) is a measure of the chemical reaction potential of a system
If ∆G is negative, the reaction is spontaneous
If ∆G is positive, the reaction is not spontaneous
Calculate the change in free energy for the following reaction
CH4(g) O2(g)  CO2(g) H2O(l)

25. Gibbs Free Energy ∆H ∆S ∆G
Enthalpy and Entropy can be combined to predict reaction spontaneity
∆G = ∆H - T∆S ∆H ∆S ∆G
Comments on Reaction - +
Always spontaneous
+ or -
Spontaneous at high temperatures
Spontaneous at low temperatures
Never spontaneous