Chapter 19 thermodynamics Laws of thermodynamics:

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1 Chapter 19 thermodynamics Laws of thermodynamics: Outline major components of the Chapter, focus of topics we covered over and over again, surprise #5, lab, demos, videos, etc. 0 th Law of Thermodynamics If A is in thermal equilibrium with B, and B is in thermal equilibrium with C, then C is also in thermal equilibrium with A. 1 st Law of Thermodynamics Energy of the universe is constant (conserved). U = q + w q = heat absorbed by the system, w = work done on the system 2 nd Law of Thermodynamics In a spontaneous process, the entropy of the universe increases. S universe = S sys + S surr > 0 (if spontaneous) For Reversible processes: S univ = S system + S surroundings = 0 For irreversible processes: S univ = S system + S surroundings > 0 3 rd Law of Thermodynamics The entropy of a pure crystalline substance at absolute zero is zero: S(0 K) = 0. Practice Exercise 19.2, p743 By considering the disorder in the reactants and products, predict whether ΔS is positive or negative for each of the following processes: a) H 2 O(l) H 2 O(g) b) Ag + (aq) + Cl - (aq) AgCl(s) c) 4 Fe(s) + 3 O 2 (g) 2 Fe 2 O 3 (s) Indicate whether each of the following reactions produces an increase or decrease in the entropy of the system: a) CO 2 (s) CO 2 (g) b) CaO(s) + CO 2 (g) CaCO 3 (s) Sample Exercise 19.3 (p. 744) The element mercury, Hg, is a silvery liquid at room temperature. The normal freezing point of mercury is o C, and its molar enthalpy of fusion is ΔH fus = 2.29 kj/mol. What is the entropy change of the system when 50.0 g of Hg(l) freezes at the normal freezing point? (ΔS sys = J/K)

2 Practice Exercise 19.3 The normal boiling point of ethanol, C 2 H 5 OH, is 78.3 o C (see Figure 11.12), and its molar enthalpy of vaporization is kj/mol. What is the change in entropy when 68.3 g of C 2 H 5 OH(g) at 1 atm pressure condenses to liquid at the normal boiling point? (-163 J/K) Sample Exercise 19.4 (p. 746) Consider the reversible melting of 1 mol of ice in a large, isothermal water bath at 0 o C and 1 atm pressure. The enthalpy of fusion of ice is 6.01 kj/mol. Calculate the entropy change in the system and in the surroundings, and the overall change in entropy of the universe for this process. (22.0 J/mol-K, J/mol-K, 0) Practice Exercise 19.4 The molar enthalpy of vaporization for liquid bromine is kj/mol. Calculate the entropy change in the system, the surroundings, and the universe for the reversible vaporization of 1 mol of liquid bromine (Br 2 ) at its normal boiling point (59 o C). (93 J/K-mol) 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 25 o C b) 2 mol of HCl(g) or 1 mol of HCl(g) at 25 o C c) 1 mol of HCl(g) or 1 mol of Ar(g) at 25 o C d) 1 mol of N 2 (s) at 24 K or 1 mol of N 2 (g) at 298 K Practice Exercise 19.5 Choose the substance with the greater entropy in each case: a) 1 mol of H 2 (g) at STP or 1 mol H 2 (g) at 100 o C and 0.5 atm b) 1 mol of H 2 O(s) at 0 o C or 1 mol of H2O(l) at 25 o C c) 1 mol of H 2 (g) at STP or 1 mol of SO 2 (g) at STP d) 1 mol of N 2 O 4 (g) at STP or 2 mol of NO 2 (g) at STP Sample Exercise 19.6 (p. 752) Predict whether the entropy change of the system in each of the following isothermal reactions is positive or negative. a) CaCO 3 (s) CaO(s) + CO 2 (g) b) N 2 (g) + 3H 2 (g) 2NH 3 (g) c) N 2 (g) + O 2 (g) 2NO(g) Practice Exercise 19.6 Predict whether ΔS is positive or negative in each of the following processes: a) HCl(g) + NH 3 (g) NH 4 Cl(s) b) 2 SO 2 (g) + O 2 (g) 2 SO 3 (g) c) Cooling nitrogen gas from 20 o C to -50 o C

3 Sample Exercise 19.7 (p. 753) Calculate ΔS for the synthesis of ammonia from N 2 (g) and H 2 (g) at 298 K. N 2 (g) + 3 H 2 (g) 2 NH 3 (g) ( J/K) Practice Exercise 19.7 Using the standard entropies in Appendix C, calculate the standard entropy change, ΔS for the following reaction at 298 K: Al 2 O 3 (s) + 3 H 2 (g) 2 Al(s) + 3 H 2 O(g) ( J/K) Practice Exercise a) Using standard enthalpies of formation and standard entropies in Appendix C, calculate ΔH and ΔS at 298 K for the following reaction: 2 SO 2 (g) + O 2 (g) 2 SO 3 (g). (ΔH = kj, ΔS = J/K) b) Using the values obtained in part (a), estimate ΔG at 400 K. (ΔG = kj) Sample Exercise (p. 762) 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, CCl 4 (l). b) What is the value of ΔGo for the equilibrium in part (a)? c) Use thermodynamic data in Appendix C and Equation to estimate the normal boiling point of CCl 4. (70 o C) Practice Exercise Use data in Appendix C to estimate the normal boiling point, in K, for elemental bromine, Br 2 (l). (The experimental value is given in Table 11.3). (330 K) Sample Exercise (p. 763) We will continue to explore the Haber process for the synthesis of ammonia: N 2 (g) + 3H 2 (g) 2 NH 3 (g) Calculate ΔG at 298 K for a reaction mixture that consists of 1.0 atm N 2, 3.0 atm H 2, and 0.50atm NH 3. (-44.9 kj/mol)

4 Practice Exercise Calculate ΔG at 298 K for the reaction of nitrogen and hydrogen to form ammonia if the reaction mixture consists of 0.50 atm N 2, 0.75 atm H 2, and 2.0 atm NH 3. (-26.0 kj/mol) Try to complete table without looking up solution. ΔH ΔS -TΔS ΔG Spontaneity Non-spontaneous Spontaneous Info on equilibrium constant, reaction quotient and Gibbs free energy za +wb yc +xd At equilibrium, Q = K eq and ΔG = 0, so: K eq = value of equilibrium constant ΔG = ΔG + RTlnQ 0 = ΔG + RTlnK eq Δ ΔG = RTlnK eq Q = reaction quotient concentration at any time in rxn From the above we can conclude: If ΔG < 0, then K eq > 1. If ΔG = 0, then K eq = 1. If ΔG > 0, then K eq < 1. Sample Exercise (p. 764) Use standard free energies of formation to calculate the equilibrium constant K eq at 25 o C for the reaction involved in the Haber process: N 2 (g) + 3 H 2 (g) 2 NH 3 (g) (7 x 10 5 ) Practice Exercise Use data from Appendix C to calculate the standard free-energy change, ΔG o, and the equilibrium constant, K eq, at 298 K for the following reaction: H 2 (g) + Br 2 (l) 2 HBr(g) ( kj/mol; 5 x ) Sample Integrative Exercise 19 (p. 766) Consider the simple salts NaCl(s) and AgCl(s). We will examine the equilibrium in which these salts dissolve in water to form aqueous solutions of ions: NaCl(s) Na + (aq) + Cl - (aq) AgCl(s) Ag + (aq) + Cl - (aq)

5 a) Calculate the value of ΔG o at 298 K for each of the preceding reactions. b) The two values from part (a) are very different. Is this difference primarily due to the enthalpy term or the entropy term of the standard free-energy change? c) Use the values of ΔG o to calculate K sp values for the two salts at 298 K. d) Sodium chloride is considered a soluble salt, whereas silver chloride is considered insoluble. Are these descriptions consistent with the answers to part (c)? e) How will ΔG o for the solution process of these salts change with increasing T? What effect should this change have on the solubility of the salts?