1. Thermodynamics Chapter 19

3. First Law of Thermodynamics You will recall from Chapter 5 that energy cannot be created or 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. © 2012 Pearson Education, Inc.

4. Spontaneous Processes and Entropy Thermodynamics lets us predict whether a process will occur but gives no information about the amount of time required for the process. A spontaneous process is one that occurs without outside intervention.

5. 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 return to vessel B. © 2012 Pearson Education, Inc.

6. Spontaneous Processes Processes that are spontaneous in one direction are nonspontaneous in the reverse direction. © 2012 Pearson Education, Inc.

7. 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. © 2012 Pearson Education, Inc.

8. Reversible v. Irreversible Processes Reversible: The universe is exactly the same as it was before the cyclic process. Irreversible: The universe is different after the cyclic process. All real processes are irreversible -- (some work is changed to heat).

9. The Second Law of Thermodynamics...in any spontaneous process there is always an increase in the entropy of the universe.  S univ > 0 for a spontaneous process.

10. Second Law of Thermodynamics In other words: For reversible processes:  S univ =  S system +  S surroundings = 0 For irreversible processes:  S univ =  S system +  S surroundings > 0 © 2012 Pearson Education, Inc.

11. Entropy The driving force for a spontaneous process is an increase in the entropy of the universe. Entropy, S, can be viewed as a measure of randomness, or disorder.

12. Entropy Like total energy, E, and enthalpy, H, entropy is a state function. Therefore,  S = S final  S initial © 2012 Pearson Education, Inc.

13. Entropy on the Molecular Scale The number of microstates and, therefore, the entropy, tends to increase with increases in –Temperature –Volume –The number of independently moving molecules. © 2012 Pearson Education, Inc.

15. Solutions Generally, when a solid is dissolved in a solvent, entropy increases. © 2012 Pearson Education, Inc.

17. Positional Entropy A gas expands into a vacuum because the expanded state has the highest positional probability of states available to the system. Therefore, S solid < S liquid << S gas

18. Entropy and Physical States Entropy increases with the freedom of motion of molecules. Therefore, S(g) > S(l) > S(s) S(g) > S(l) > S(s) © 2012 Pearson Education, Inc.

20. Third Law of Thermodynamics The entropy of a pure crystalline substance at absolute zero is 0. © 2012 Pearson Education, Inc.

21. Entropy

22. Standard Entropies These are molar entropy values of substances in their standard states. Standard entropies tend to increase with increasing molar mass. © 2012 Pearson Education, Inc.

23. Standard Entropies Larger and more complex molecules have greater entropies. © 2012 Pearson Education, Inc.

24. Entropy Changes Entropy changes for a reaction can be estimated in a manner analogous to that by which  H is estimated:  S  =  n  S  (products) —  m  S  (reactants) where n and m are the coefficients in the balanced chemical equation. © 2012 Pearson Education, Inc.

25. Entropy Changes in Surroundings Heat that flows into or out of the system changes the entropy of the surroundings. For an isothermal process:  S surr =  q sys T At constant pressure, q sys is simply  H  for the system. © 2012 Pearson Education, Inc.

26. Entropy Change in the Universe The universe is composed of the system and the surroundings. Therefore,  S universe =  S system +  S surroundings For spontaneous processes  S universe > 0  S universe > 0 © 2012 Pearson Education, Inc.

27. Free Energy  G =  H  T  S (from the standpoint of the system) A process (at constant T, P) is spontaneous in the direction in which free energy decreases:  G means +  S univ

28. Gibbs Free Energy 1.If  G is negative, the forward reaction is spontaneous. 2.If  G is 0, the system is at equilibrium. 3.If  G is positive, the reaction is spontaneous in the reverse direction. © 2012 Pearson Education, Inc.

29. 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 then comes from the entropy term. © 2012 Pearson Education, Inc.

30. Effect of  H and  S on Spontaneity

32. Free Energy Change and Chemical Reactions  G  = standard free energy change that occurs if reactants in their standard state are converted to products in their standard state.  G  =  n p  G f  (products)   n r  G f  (reactants)

33. Free Energy and Equilibrium Under any conditions, standard or nonstandard, the free energy change can be found this way:  G =  G  + RT ln Q (Under standard conditions, all concentrations are 1 M, so Q = 1 and ln Q = 0; the last term drops out.) © 2012 Pearson Education, Inc.

34. Free Energy and Pressure  G =  G  + RT ln(Q) Q = reaction quotient from the law of mass action.

35. Free Energy and Equilibrium  G  =  RT ln(K) K = equilibrium constant This is so because  G = 0 and Q = K at equilibrium.

39. Temperature Dependence of K y = mx + b (  H  and S   independent of temperature over a small temperature range) ln(K) = -  H o /R*(1/T) +  S o /R

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