1. Chapter 17 Free Energy and Thermodynamics

2. First law of thermodynamics: Energy cannot be created or destroyed.
The total energy of the universe cannot change.
DEuniverse = 0 = DEsystem + DEsurroundings

3. two ways energy moves in and out of systems:
For an exothermic reaction, “lost” heat from the system goes into the surroundings.
For an endothermic reaction, “gained” heat goes from the surroundings into the system.
two ways energy moves in and out of systems:
As heat, q
As work, w

4. Thermodynamics
Thermodynamics predicts whether a process will proceed (occur) under the given conditions.
= spontaneous process
Spontaneity is determined by comparing the chemical potential energy of the system before the reaction with the free energy of the system after the reaction.
spontaneity ≠ fast or slow

5. Spontaneous processes are 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.

6. Thermodynamics vs. Kinetics

7. There are two factors that determine the spontaneity of a change:
The enthalpy change, DH, is the difference in the sum of the internal energy and PV work energy of the reactants compared to the products.
The entropy change, DS, is the difference in randomness of the reactants compared to the products.

8. Enthalpy
The heat energy change at constant pressure for a physical or chemical change.
DH is generally measured in kJ/mol.
exothermic = energy released, DH is negative
endothermic = energy absorbed, DH is positive
The enthalpy is favorable for exothermic reactions and unfavorable for endothermic reactions.

9. Entropy
Entropy, S, is a thermodynamic function that increases as the number of energetically equivalent ways of arranging the components increases. (CHAOS)
S is generally measured in J/mol.
Increasing entropy,  S is positive
Decreasing entropy, ∆ S is negative
Increasing entropy is favorable for reactions

10. Changes in Entropy, DS Some changes that increase the entropy are
reactions whose products are in a more random state (like gas phase rather than solid)
reactions that have larger numbers of product molecules than reactant molecules
increase in temperature
solids dissociating into ions upon dissolving

11. Practice—Predict whether DSsystem is + or − for each of the following.
Heating air in a balloon
Water vapor condensing
Separation of oil and vinegar salad dressing
Dissolving sugar in tea
2 HgO(s)  2 Hg(l) + O2(g)
2 NH3(g)  N2(g) + 3 H2(g)
Ag+(aq) + Cl−(aq)  AgCl(s)

12. The Second Law of Thermodynamics
The total entropy change of the universe must be positive (increase) for a process to be spontaneous.
DSuniverse = DSsystem + Dssurroundings
If the entropy of the system decreases, then the entropy of the surroundings must increase by a larger amount.
When DSsystem is negative, DSsurroundings is positive.

13. UIL: Temperature Dependence of DSsurroundings
When a system process is exothermic, it adds heat to the surroundings, increasing the entropy of the surroundings.
When a system process is endothermic, it takes heat from the surroundings, decreasing the entropy of the surroundings.
The amount the entropy of the surroundings changes depends on its initial temperature.
The higher the original temperature, the less effect addition or removal of heat has.

14. Calculating DH and DS
∆Hsys = ∆Hrxn = Σ(H°prd) − Σ(H°rct)

15. The Third Law of Thermodynamics— Absolute Entropy
The absolute entropy of a substance is the amount of energy it has due to dispersion of energy through its particles.
The third law states that for a perfect crystal at absolute zero, the absolute entropy = 0 J/mol∙K.
Therefore, every substance that is not a perfect crystal at absolute zero has some energy from entropy.
Therefore, the absolute entropy of substances is always +.

16. Standard Absolute Entropies
entropies for 1 mole of a substance at 298 K
1) state (gas>liquid>solid)
2) allotrope form (more open structures have greater entropy)
3) molecular complexity and mass (more = greater entropy)
4) degree of dissolution (aq > entropy)

17. Calculate DS for the reaction 2 H2(g) + O2(g)  2 H2O(g)

18. Calculate DS for the reaction 4 NH3(g) + 5 O2(g)  4 NO(g) + 6 H2O(g)

19. Gibbs Free Energy and Spontaneity
The Gibbs free energy is the maximum amount of work energy that can be released to the surroundings by a system.
The Gibbs free energy is often called the chemical potential energy because it is analogous to the storing of energy in a mechanical system.
DGsys = DHsys−TDSsys
Spontaneous changes have negative free energy change. (∆G is negative)

20. If ∆G = 0, the system is at equilibrium.

21. The reaction CCl4(g)  C(s, graphite) + 2 Cl2(g) has
DH = kJ and DS = J/K at 25 °C. Calculate DG and determine whether it is spontaneous.

22. The reaction Al(s) + Fe2O3(s)  Fe(s) + Al2O3(s) has
DH = −847.6 kJ and DS = −41.3 J/K at 25 °C. Calculate DG and determine whether it is spontaneous.

23. The reaction CCl4(g)  C(s, graphite) + 2 Cl2(g)
has DH = kJ and DS = J/K. Calculate the minimum temperature for it to be spontaneous.

24. DGorxn = SnGof (prd) − SnGof (rct)
at 25 C:
DGorxn = SnGof (prd) − SnGof (rct)
Determine the free energy
change in the following
reaction at 298 K:
2 H2O(g) + O2(g)  2 H2O2(g).
DH, kJ/mol
S, J/mol
H2O2(g)
−136.3
232.7
O2(g)
205.2
H2O(g)
−241.8
188.8

25. Practice—Determine the free energy change in the following reaction at 298 K: 2 H2O(g) + O2(g)  2 H2O2(g)
Given:
H2(g) + O2(g)  H2O2(g) DGº = −105.6 kJ
2 H2(g) + O2(g)  2 H2O(g) DGº = −457.2 kJ

26. UIL needs this: DG under Nonstandard Conditions
DG = DG only when the reactants and products are in their standard states.
their normal state at that temperature
partial pressure of gas = 1 atm
concentration = 1 M
Under nonstandard conditions,
DG = DG + RTlnQ (Q is the reaction quotient)
At equilibrium, DG = 0 = −RTlnK