1. Chapter 19 Chemical Thermodynamics John D. Bookstaver St. Charles Community College St. Peters, MO 2006, Prentice Hall, Inc. Modified by S.A. Green, 2006 Modified by D. Amuso 2011

2. Thermodynamics The study of the relationships between heat, work, and the energy of a system.

3. First Law of Thermodynamics You will recall from Chapter 5 that energy cannot be created nor destroyed. Therefore, the total energy of the universe is a constant. Energy can, however, be converted from one form to another or transferred from a system to the surroundings or vice versa.

4. Second Law of Thermodynamics The total Entropy of the universe increases in any spontaneous, irreversible process.

5. Entropy Entropy can be thought of as a measure of the randomness or disorder of a system. It is related to the various modes of motion in molecules (microstates). Vibrational Rotational Translational

6. Entropy Molecules have more entropy (disorder) when: 1)Phase Changes from: S  L  G Example: Sublimation CO 2 (s)  CO 2 (g) 2)Dissolving occurs (solution forms): Example: NaCl(s)  Na + (aq) + Cl - (aq)

7. Entropy 3) Temperature increases Example: Fe(s) at 0 o C  Fe(s) at 25 o C 4) For Gases ONLY, when Volume increases or Pressure decreases Examples: 2 Liters He(g)  4 Liters He(g) 3 atm He(g)  1 atm He(g)

8. Entropy 5) Rx results in more molecules/moles of gas Examples: 2NH 3 (g)  N 2 (g) + 3H 2 (g) CaCO 3 (s)  CaO(s) + CO 2 (g) N 2 O 4 (g)  2 NO 2 (g) This one is difficult to predict: N 2 (g) + O 2 (g)  2 NO(g)

9. Entropy 6)When there are more moles Example: 1 mole H 2 O(g)  2 moles H 2 O(g) 7) When there are more atoms per molecule Examples: 1 mole Ar(g)  1 mole HCl(g) 1 mole NO 2 (g)  1 mole N 2 O 4 (g)

10. Entropy 8) When an atom has a bigger atomic number 1 mole He(g)  1 mole Ne (g)

11. Spontaneous Processes Spontaneous processes are those that can proceed without any outside intervention. The gas in vessel B will spontaneously effuse into vessel A, but once the gas is in both vessels, it will not spontaneously move to just one vessel.

12. Spontaneous Processes Processes that are spontaneous at one temperature may be nonspontaneous at other temperatures. Above 0  C it is spontaneous for ice to melt. Below 0  C the reverse process is spontaneous.

13. Spontaneous Processes Processes that are spontaneous in one direction are nonspontaneous in the reverse direction. Spontaneous processes are irreversible.

14. Second Law of Thermodynamics A reversible process results in no overall change in Entropy while an irreversible, spontaneous process results in an overall increase in Entropy. Reversible: Irreversible (Spontaneous):

15. Third Law of Thermodynamics The Entropy of a pure crystalline substance at absolute zero is zero. why?

16. Third Law of Thermodynamics

17. Standard Entropies Standard Conditions: 298 K, 1 atm, 1 Molar The values for Standard Entropies (S o ) are expressed in J/mol-K. Note: Increase with increasing molar mass.

18. Standard Entropies Larger and more complex molecules have greater entropies.

19. Entropy Changes  S o =  n S o products -  m  S o reactants Be careful: S°units are in J/mol-K Note for pure elements:

20. Gibbs Free Energy Use  G to decide if a process is spontaneous  G = negative value = spontaneous  G = zero = at equilibrium  G = positive value = not spontaneous Note: equation can be used w/o the o too.  G o =  H o – T  S o

21. Gibbs Free Energy 1.If  G is negative  G = maximum amt of energy ‘free’ to do work by the reaction 2. If  G is positive  G = minimum amt of work needed to make the reaction happen  G o =  H o – T  S o

22. Gibbs Free Energy In our tables, units are:  G o = kJ/mol  H o = kJ/mol  S o = J/mol-K  G o =  H o – T  S o

23. Free Energy and Temperature There are two parts to the free energy equation:   H  — the enthalpy term  T  S  — the entropy term The temperature dependence of free energy comes from the entropy term.

24. What causes a reaction to be spontaneous? Think Humpty Dumpty System tend to seek:  Minimum Enthalpy Exothermic Rx,  H = negative  Maximum Entropy More disorder,  S = positive Because:  G o =  H o – T  S o - = (-) - (+)

25. Free Energy and Temperature By knowing the sign (+ or -) of  S and  H, we can get the sign of  G and determine if a reaction is spontaneous.

26. At Equilibrium  G o = zero Therefore:  H o = T  S o Or:  S o  H o T Use this equation when asked to calculate enthalpy of vaporization or enthalpy of fusion.  G o =  H o – T  S o

27.  G o =  n G o f products -  m  G o f reactants Standard Free Energy Changes Be careful: Values for  G f  are in kJ/mol  G  can be looked up in tables or calculated from S° and  H .   G o =  H o – T  S o

28. Free Energy and Equilibrium Remember from above: If  G is 0, the system is at equilibrium. So  G must be related to the equilibrium constant, K (chapter 15). The standard free energy,  G°, is directly linked to K eq by:  G o =  RT ln K   Where R = 8.314 J/mol-K

29.  G o =  RT ln K   If the free energy change is a negative value,  the reaction is spontaneous, ln K must be a positive value, and K will be a large number meaning the equilibrium mixture is mainly products.  If the free energy change is zero,  ln K = zero and K = one. Relationships

30. Free Energy and Equilibrium Under non-standard conditions, we need to use  G instead of  G°. Q is the reaction quotient from chapter 15.