1. Chemical Thermodynamics First Law of Thermodynamics You will recall from earlier this year 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.

2. Chemical Thermodynamics Copyright © Cengage Learning. All rights reserved 2 Thermodynamics vs. Kinetics Domain of Kinetics  Rate of a reaction depends on the pathway from reactants to products. Thermodynamics tells us whether a reaction is spontaneous based only on the properties of reactants and products.

3. Chemical Thermodynamics © 2009, Prentice-Hall, Inc. Spontaneous Processes Spontaneous processes are those that can proceed without any outside intervention.

4. Chemical Thermodynamics Spontaneous Processes Processes that are spontaneous in one direction are nonspontaneous in the reverse direction.

5. Chemical Thermodynamics 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.

6. Chemical Thermodynamics Entropy (S) Entropy can be thought of as a measure of the randomness or disorder of a system. Nature favors changes which result in higher entropy

7. Chemical Thermodynamics Positional Entropy The number of possible arrangements of molecules. (microstates) If 4 molecules were placed inside the two-bulbed container, how many arrangements can be achieved?

8. Chemical Thermodynamics

9. Chemical Thermodynamics Entropy on the Molecular Scale Molecules exhibit several types of motion:  Translational: Movement of the entire molecule from one place to another.  Vibrational: Periodic motion of atoms within a molecule.  Rotational: Rotation of the molecule on about an axis or rotation about  bonds.

10. Chemical Thermodynamics Entropy Changes In general, entropy increases when  Gases are formed from liquids and solids.  Liquids or solutions are formed from solids.  The number of gas molecules increases.  The number of moles increases.  The temperature increases.

11. Chemical Thermodynamics Like total energy, E, and enthalpy, H, entropy is a state function. Therefore,  S = S final  S initial Predict the sign of ΔS for the following: a)The evaporation of alcohol b)The freezing of water c)Compressing an ideal gas at constant temperature d)Heating an ideal gas at constant pressure e)Dissolving NaCl in water

12. Chemical Thermodynamics Second Law of Thermodynamics In any spontaneous process there is always an increase in the entropy of the universe. The total energy of the universe is constant, but the entropy is increasing. ΔS surr = +; entropy of the universe increases ΔS surr = -; process is spontaneous in opposite direction ΔS surr = 0; process has no tendency to occur

13. Chemical Thermodynamics The sign of ΔS surr depends on the direction of the heat flow. H 2 O(l)  H 2 O(g) Is this reaction endothermic or exothermic? What happens to Δs surr ? Why are exothermic reactions favored? (Why does nature want to be “lazy”?)

14. Chemical Thermodynamics The magnitude of ΔS surr depends on the temperature. Δs surr depends directly on the amount of heat transferred and inversely on the temperature The minus sign is needed because ΔH is determined with respect to the system and this equation expresses ΔS of the surroundings

15. Chemical Thermodynamics Calculate ΔS surr Each reaction occurs at 25°C and 1 atm H 2 O(l)  H 2 O(g)ΔH = +44 kJ C 3 H 8 (g) + 5O 2 (g)  3CO 2 (g) + 4H 2 O(g) ΔH = -2045 kJ

16. Chemical Thermodynamics

17. Chemical Thermodynamics Third Law of Thermodynamics The entropy of a pure crystalline substance at absolute zero is 0.

18. Chemical Thermodynamics These are molar entropy values of substances in their standard states. Standard entropies tend to increase with increasing molar mass. Standard entropy value ΔS °

19. Chemical Thermodynamics Standard Entropies Larger and more complex molecules have greater entropies.

20. Chemical Thermodynamics Entropy Changes Entropy changes for a reaction can be calculated the same way we used for  H: S° for each component is found in a table. Note for pure elements:

21. Chemical Thermodynamics Gibbs Free Energy Substituting into Yields: Dividing by T:

22. Chemical Thermodynamics Gibbs Free Energy Negative ΔG means positive Δs univ (which means spontaneous)

23. Chemical Thermodynamics 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.

24. Chemical Thermodynamics Calculate Free Energy H 2 O(s)  H 2 O(l) ΔH ° = 6060 J/mol ΔS ° = 22.1 J/K ·mol At 10 °C At -10 °C

25. Chemical Thermodynamics 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. Chemical Thermodynamics Br 2 (l)  Br 2 (g) ΔH ° = 31.0 kJ/mol ΔS ° = 93.0 J/K ·mol What is the normal boiling point of bromine?

27. Chemical Thermodynamics Standard Free Energy Changes Standard free energies of formation,  G f  are analogous to standard enthalpies of formation,  H f .  G  can be looked up in tables, or calculated from S° and  H .

28. Chemical Thermodynamics Free Energy and Equilibrium  G = 0, K eq = 1 (the system is at equilibrium)  G 1 (reaction proceeds  )  G > 0, K eq < 1 (reaction proceeds  )

29. Chemical Thermodynamics Free Energy and Work In addition to predicting the spontaneity of reactions, free energy can tell us how much work can be done by a process w max = ΔG

30. Chemical Thermodynamics Achieving the maximum work available from a spontaneous process can occur only via a hypothetical pathway. Any real pathway wastes energy. All real processes are irreversible. First law: You can’t win, you can only break even. Second law: You can’t break even.