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Chemical Thermodynamics

Chemical Thermodynamics. Thermo Part 2. 1. 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.

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Chemical Thermodynamics

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  1. Chemical Thermodynamics Thermo Part 2 1

  2. 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. 2

  3. 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 go back to one container 3

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

  5. 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. 5

  6. SAMPLE EXERCISE 19.1 Identifying Spontaneous Processes Predict whether the following processes are spontaneous as described, spontaneous in the reverse direction, or in equilibrium: (a) When a piece of metal heated to 150ºC is added to water at 40ºC, the water gets hotter. (b) Water at room temperature decomposes into H2(g) and O2(g). (c) Benzene vapor, C6H6(g), at a pressure of 1 atm condenses to liquid benzene at the normal boiling point of benzene, 80.1ºC. PRACTICE EXERCISE Under 1 atm pressure CO2(s) sublimes at –78ºC. Is the transformation of CO2(s) to CO2(g) a spontaneous process at –100ºC and 1 atm pressure? 6

  7. Reversible Processes In a reversible process the system changes in such a way that the system and surroundings can be put back in their original states by exactly reversing the process. 7

  8. Irreversible Processes • Irreversible processes cannot be undone by exactly reversing the change to the system. • Spontaneous processes are irreversible. 8

  9. q T Entropy • Entropy (S) is a term coined by Rudolph Clausius in the 19th century. • Clausius was convinced of the significance of the ratio of heat delivered and the temperature at which it is delivered, • Entropy can be thought of as a measure of the randomness of a system. • It is related to the various modes of motion in molecules. 9

  10. Entropy • Like total energy, E, and enthalpy, H, entropy is a state function. • Therefore, S = SfinalSinitial 10

  11. qrev T S = Entropy • For a process occurring at constant temperature (an isothermal process), the change in entropy is equal to the heat that would be transferred if the process were reversible divided by the temperature: 11

  12. Second Law of Thermodynamics The second law of thermodynamics states that the entropy of the universe increases for spontaneous processes, and the entropy of the universe does not change for reversible processes. 12

  13. Second Law of Thermodynamics In other words: For reversible processes: Suniv = Ssystem + Ssurroundings = 0 For irreversible processes: Suniv = Ssystem + Ssurroundings > 0 S is entropy. H is enthalpy. 13

  14. Second Law of Thermodynamics These last truths mean that as a result of all spontaneous processes the entropy of the universe increases. 14

  15. Entropy on the Molecular Scale • Ludwig Boltzmann described the concept of entropy on the molecular level. • Temperature is a measure of the average kinetic energy of the molecules in a sample. 15

  16. 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. 16

  17. Entropy on the Molecular Scale • Boltzmann envisioned the motions of a sample of molecules at a particular instant in time. • This would be akin to taking a snapshot of all the molecules. • He referred to this sampling as a microstate of the thermodynamic system. 17

  18. Entropy on the Molecular Scale • Each thermodynamic state has a specific number of microstates, W, associated with it. • Entropy is S = k lnW where k is the Boltzmann constant, 1.38  1023 J/K. 18

  19. S= k ln lnWfinal lnWinitial Entropy on the Molecular Scale • The change in entropy for a process, then, is S = k lnWfinalk lnWinitial • Entropy increases with the number of microstates in the system. 19

  20. 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. 20

  21. Entropy and Physical States • Entropy increases with the freedom of motion of molecules. • Therefore, S(g) > S(l) > S(s) 21

  22. Solutions Generally, when a solid is dissolved in a solvent, entropy increases. 22

  23. 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. 23

  24. SAMPLE EXERCISE 19.3 Predicting the Sign of S Predict whether S is positive or negative for each of the following processes, assuming each occurs at constant temperature: 24

  25. PRACTICE EXERCISE Indicate whether each of the following processes produces an increase or decrease in the entropy of the system: 25

  26. SAMPLE EXERCISE 19.4 Predicting Which Sample of Matter Has the Higher Entropy Choose the sample of matter that has greater entropy in each pair, and explain your choice: (a) 1 mol of NaCl(s) or 1 mol of HCl(g) at 25ºC, (b) 2 mol of HCl(g) or 1 mol of HCl(g) at 25ºC, (c) 1 mol of HCl(g) or 1 mol of Ar(g) at 298 K. PRACTICE EXERCISE Choose the substance with the greater entropy in each case: (a) 1 mol of H2(g) at STP or 1 mol of H2(g) at 100ºC and 0.5 atm, (b) 1 mol of H2O(s) at 0ºC or 1 mol of H2O(l) at 25ºC, (c) 1 mol of H2(g) at STP or 1 mol of SO2(g) at STP, (d) 1 mol of N2O4(g) at STP or 2 mol of NO2(g) at STP. 26

  27. Third Law of Thermodynamics The entropy of a pure crystalline substance at absolute zero is 0. 27

  28. Standard Entropies • These are molar entropy values of substances in their standard states. • Standard entropies tend to increase with increasing molar mass. 28

  29. Standard Entropies Larger and more complex molecules have greater entropies. (More stuff, more entropy) 29

  30. 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. 30

  31. PRACTICE EXERCISE Using the standard entropies in Appendix C, calculate the standard entropy change, S°, for the following reaction at 298 K: SAMPLE EXERCISE 19.5 Calculating S from Tabulated Entropies Calculate S° for the synthesis of ammonia from N2(g) and H2(g) at 298 K: Step 1 Formula with substitutions: Step 2 Gather data from chart or question provided. We would fine these in Appendix C of your book: 31

  32. qsys T Ssurr = Entropy Changes in Surroundings • Heat that flows into or out of the system changes the entropy of the surroundings. • For an isothermal process: • At constant pressure, qsys is simply H for the system. 32

  33. Entropy Change in the Universe • The universe is composed of the system and the surroundings. • Therefore, Suniverse = Ssystem + Ssurroundings • For spontaneous processes Suniverse > 0 33

  34. Hsystem T Entropy Change in the Universe • This becomes: Suniverse = Ssystem + Multiplying both sides by T, TSuniverse = Hsystem  TSsystem 34

  35. Gibbs Free Energy • TDSuniverse is defined as the Gibbs free energy, G. • When Suniverse is positive, G is negative. • Therefore, when G is negative, a process is spontaneous. 35

  36. Gibbs Free Energy • If DG is negative, the forward reaction is spontaneous. • If DG is 0, the system is at equilibrium. • If G is positive, the reaction is spontaneous in the reverse direction. 36

  37. DG = SnDG(products)  SmG(reactants) f f Standard Free Energy Changes Analogous to standard enthalpies of formation are standard free energies of formation, G. f where n and m are the stoichiometric coefficients. 37

  38. (a) By using data from Appendix C, calculate the standard free-energy change for the following reaction at 298 K: SAMPLE EXERCISE 19.6 Calculating Standard Free-Energy Change from Free Energies of Formation (b) What is Gºfor the reverse of the above reaction? 38

  39. PRACTICE EXERCISE By using data from Appendix C, calculate Gº at 298 K for the combustion of methane: 39

  40. In Section 5.7 we used Hess’s law to calculate Hº for the combustion of propane gas at 298 K: SAMPLE EXERCISE 19.7 Estimating and Calculating G° (a)Without using data from Appendix C, predict whether Gº for this reaction is more negative or less negative than Hº. (b) Use data from Appendix C to calculate the standard free-energy change for the reaction at 298 K. Is your prediction from part (a) correct? 40

  41. PRACTICE EXERCISE Consider the combustion of propane to form CO2(g) and H2O(g) at 298 K: Would you expect Gº to be more negative or less negative than H°? 41

  42. Free Energy Changes At temperatures other than 25°C, DG° = DH  TS How does G change with temperature? 42

  43. 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. 43

  44. Free Energy and Temperature 44

  45. The Haber process for the production of ammonia involves the equilibrium SAMPLE EXERCISE 19.8 Determining the Effect of Temperature on Spontaneity Assume that Hº and Sº for this reaction do not change with temperature. (a) Predict the direction in which Gº for this reaction changes with increasing temperature. (b) Calculate the values of Gº for the reaction at 25ºC and 500ºC. 45

  46. 298 K for the following reaction: (b) Using the values obtained in part (a), estimate Gº at 400 K. SAMPLE EXERCISE 19.8continued PRACTICE EXERCISE (a) Using standard enthalpies of formation and standard entropies in Appendix C, calculate Hº and Sº at 46

  47. Free Energy and Equilibrium Under any conditions, standard or nonstandard, the free energy change can be found this way: G = G + RT lnQ (Under standard conditions, all concentrations are 1 M, so Q = 1 and lnQ = 0; the last term drops out.) 47

  48. SAMPLE EXERCISE 19.9 Relating G to a Phase Change at Equilibrium As we saw in Section 11.5, the normal boiling point is the temperature at which a pure liquid is in equilibrium with its vapor at a pressure of 1 atm. (a) Write the chemical equation that defines the normal boiling point of liquid carbon tetrachloride, CCl4(l). (b) What is the value of Gº for the equilibrium in part (a)? (c) Use thermodynamic data in Appendix C and Equation 19.20 to estimate the normal boiling point of CCl4. 48

  49. PRACTICE EXERCISE Use data in Appendix C to estimate the normal boiling point, in K, for elemental bromine, Br2(l). (The experimental value is given in Table 11.3.) 49

  50. We will continue to explore the Haber process for the synthesis of ammonia: Calculate G at 298 K for a reaction mixture that consists of 1.0 atm N2, 3.0 atm H2, and 0.50 atm NH3. SAMPLE EXERCISE 19.10 Calculating the Free-Energy Change Under Nonstandard Conditions 50

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