Overall Objectives – GC II

1. Overall Objectives – GC II • Thermodynamics: Directionality of Chemical Reactions (Chapter 18) • Liquids, Solids, and Materials (Chapter 11) • Solutes and Solutions (Chapter 15) • Chemical Kinetics: Rates of Reactions (Chapter 13) • Chemical Equilibrium (Chapter 14) • Acids and Bases (Chapter 16) • Additional Aqueous Equilibria (Chapter 17) • Electrochemistry (Chapter 19) All chapters share themes of thermodynamics and equilibria.

2. Main ‘Thread’– GC II • Thermodynamics: Directionality of Chemical Reactions (Chapter 18) • Chemical Kinetics: Rates of Reactions (Chapter 13) • Chemical Equilibrium (Chapter 14) Will a reaction proceed? How fast will it go? How far will it go?

3. Chemistry: The Molecular Science Moore, Stanitski, and Jurs Chapter 18: Thermodynamics: Directionality of Chemical Reactions Chemistry: The Science in Context Gilbert, Kriss, Davies Chapter 13: Entropy and Free Energy Chemistry: Principles and Reactions Masterton & Hurley Chapter 17: Spontaneous Processes

4. Thermodynamics: Directionality of Chemical Reactions (Chapter 18) Will a reaction proceed? • Introduction (18.1, 18.2) • What is a spontaneous reaction? • What is entropy (ΔS)? • Qualitative Guidelines for Entropy, ΔS (18.3) • Entropy and Criteria for spontaneity (18.5) • Second Law of thermodynamics • Gibbs Free Energy, ΔG (18.6) • The ultimate criteria for spontaneity • Free Energy ΔG and Equilibria Constants (18.7)

5. Spontaneous? • What is a spontaneous process? A spontaneous reaction will occur “by itself” , at a particular temperature and pressure, without exertion of any outside forces. • Which of these reactions are spontaneous (at 1atm and 25°C)? • 1. Reaction of sodium metal in water • Na (s) + H2O (l) Na+(aq) + OH- (aq) + H2(g) • 2. Iron rusting • 4Fe (s) + 3O2(g) + 6H2O (l) 4Fe(OH)3 (s) • 3. Mixing of two ideal gases (helium and argon) • 4. Melting of ice in your soda on a hot day • H2O (s) H2O (l) SLOW!

6. The Energy Factor: Enthalpy, ΔHMany spontaneous processes proceed with a decrease of energy: exothermic reaction, net bond-making, more stable products, “down-hill”. 1. Reaction of sodium metal in water Na (s) + H2O (l)  Na+(aq) + OH-(aq) + H2 (g) 2. Rusting 4Fe (s) + 3O2(g) + 6H2O (l) 4Fe(OH)3 (s) 3. Mixing of two ideal gases (He + Ar) 4. Melting of ice in your soda on a hot day H2O (s)  H2O (l) (ΔH< O, exothermic) (ΔH< O, exothermic) (ΔH= O) (ΔH> O, endothermic)

7. The Randomness Factor: Entropy (∆S) Nature tends to move spontaneously from a state of lower probability to one of higher probability, or ordered to more random. Second Law of thermodynamics A process occurs spontaneously if it results in an INCREASE in the ENTROPY of the UNIVERSE!!! Entropy (S)is a measure of nanoscale disorder. • dispersal of energy • dispersal of matter probability driven

8. Entropy (∆S): Dispersal of Energy Exothermic reactions: • E is transferred to the surroundings. • Bond (potential) E spread over many particles rather than concentrated over a few. • E is distributed more randomly.

9. Entropy (∆S): Dispersal of Energy Consider 2 atoms (A and B) and 2 packets of energy (* *). Possible distributions of energy: A** or A*,B* or B** Now consider contact with atoms C and D: How are the 2 packets of energy dispersed? A** B C D A* B* C D A* B C* D A* B C D* A B** C D A B* C* D A B* C D* A B C** D A B C* D* A B C D** Only 3/10 (30%) have the energy in atoms A and B like the original state. There is 70% chance that E will spread out over more atoms. E concentration is much less likely.

10. Entropy (∆S): Dispersal of Matter Matter tends to disperse (like E ): • Matter is dispersed: • Pure systems mix. • Particles occupy a larger space. Gases expand (if no barrier). 50% chance that a molecule will be in a given flask. ForN molecules the chance that ALL will be in the top flask = (½)N (½)2 = 0.25 (½)1000 = 8 x 10-31

11. Entropy (∆S): Dispersal of Matter • Matter is dispersed: • Pure systems mix. • Particles occupy a larger space. dice marbles CLOSED OPEN ordered random SPONTANEOUS

12. Qualitative Prediction Of ΔSsystem Entropy usually increases in the following situations: • gas >> liquid > solid (of the same substance) • temperature increases (more motion) • number of moles increases (ReactantsProducts) • total volume increases • solid or liquid substance dissolving

13. Qualitative Prediction Of ΔSsystem Sgas >> Sliquid > Ssolid more freedom = more disorder = higher S gas liquid solid S°(I2) 261 - 116 S°(Br2) 245 152 - S°(Cl2) 223 - - room temperature Weaker bond = larger S (Less tightly bound = easier motion)

14. Qualitative Guidelines for Entropy Solid or liquid dissolving Larger matter dispersal. S usually* increases: Substance S° (pure) S° (aq) CH3COOH(l) 160 179 NH4NO3(s) 151 260

15. What Happens When Gases Dissolve? Gas-molecule motion becomes restricted. ΔS < 0.

16. Entropy and Temperature A perfect crystal at 0K has an entropy of 0. 0K

17. Predicting Entropy Changes Will entropy increase or decrease for the following processes? Increase entropy if…. number of moles increases (ReactantsProducts) 2 CO(g) + O2(g) → 2 CO2(g) Decrease 3 non-identical gas molecules →2 identical gas molecules. Less disorder. NaCl(s) → Na+(aq) + Cl-(aq) Each crystal “unit” → 2 ions dispersed in water. Much more disorder. Increase H2O(l) → H2O(s) Decrease Disorganized water molecules → more organized ice molecules. Less disorder.

18. Sign of ΔSsystem Predict (qualitatively) the sign of ΔSsystem for the following: CaCO3(s)  CaO (s) + CO2 (g) CS2(l)  CS2(g) 2 Hg(l) + O2(g)  2 HgO (s) 2 Na2O2(s) + 2 H2O (l)  4 NaOH (aq) + O2(g) + + - +

19. (ΔSsystem) alone good enough? 1. Melting of ice in your soda on a hot day H2O (s)  H2O (l) Spontaneous and (ΔSsystem > 0) 2. Freezing of water on a cold day (-15oC) H2O (l)  H2O (s) Spontaneous and (ΔSsystem < 0)

20. Gibbs Free Energy Neither entropy (S), nor enthalpy (H ), alone can predict whether a reaction is spontaneous. Spontaneous reactions can: - be exothermic or endothermic. -increase or decrease the Ssystem. The Gibbs free energy (G), combines H and S. ΔGcan predict if a reaction is product favored or not.

21. Gibbs Free Energy For a constant T, ΔG is equal to: ΔG = ΔH – TΔS(T in K). If G: Decreases(ΔG < 0) a reaction is spontaneous. Increases(ΔG > 0) a reaction is nonspontaneous. (at constant P and T).

22. Free Energy of System (∆Gsystem) ∆Gsystem = (∆Hsystem - T ∆Ssystem ) < 0 Ultimate Criteria for Spontaneity If ∆Gsystem < 0 (negative) process is spontaneous or exergonic If ∆Gsystem > 0 (positive) process is not spontaneous or endergonic If ∆Gsystem = 0 process is in equilibrium

23. Spontaneity - Driving Forces • Thermodynamic Criterea for Reaction Spontaneity • ΔH, the change in enthalpy ΔH < 0 ΔH < 0, tends to make a reaction spontaneous • ΔS, the change in entropy ΔS > 0, tends to make a reaction spontaneous • ΔG, the change in free energy ΔG < 0, will be spontaneous .

24. The Effect of T on Reaction Direction Since ΔG = ΔH – TΔS Sign of ΔH Sign of ΔS Product Favored? Negative (exothermic) Positive Yes Negative (exothermic) Negative Yes at low T; No at high T Positive (endothermic) Positive No at low T;Yes at high T Positive (endothermic) Negative No ΔG = ΔH – TΔS

25. Putting it all together (∆G = ∆H- T ∆S) qualitative prediction Classify reaction into 4 types : (always spontaneous, never spontaneous, spontaneous high T, spontaneous at low T) a) 2 O3 (g)  3 O2(g) ∆H< 0, ∆S = ∆G = b) 2 C(graphite)+ 2 H2 (g)  C2H4 (g) ∆H> 0, ∆S = ∆G = c) N2(g) + 3H2(g)  2 NH3(g) ∆H< 0, ∆S = ∆G = d) 2 H2O (g)  2 H2 (g) + O2 (g) ∆H> 0, ∆S = ∆G = + - - + - + - low T - high T DEMO: rubber band

26. The Effect of T on Reaction Direction ΔG = ΔH − TΔS Reactions go from non-spontaneous (+ΔG) to spontaneous (-ΔG) when, TΔS > ΔH ∆H +, ∆S+ Need to maximize Entropy term. T> ΔH ΔS

27. The Effect of T on Reaction Direction At what T will the following reaction be spontaneous? CH4(g) + H2O(g) CO(g) + 3 H2(g) ΔH° = 206.10 kJ ΔS° = 214.63 J K-1 ΔG = ΔH − TΔS Spontaneous at higher T. Switch occurs at ΔG = 0, when two opposing terms are equal 0= ΔH − TΔS T= ΔH/ΔS T = 206.10 kJ/mol = 960 K 0.21463 kJ/Kmol Spontaneous > 960K

28. The Effect of T on Reaction Direction • Do Now: • Calculate the temperature at which the following reaction will occur spontaneously: Fe2O3(s) +3H2(g) 2Fe(s) +3H2O(g) • At what T will the following reaction be spontaneous? H2O (l) H2O (g) ΔH° = 98.8 kJ ΔS° = 141.5 J K-1 698K 373K, 100°C ΔH° = 40.7 kJ ΔS° = 109 J K-1

29. Reactions that Reach Equilibrium ? ΔG° = −RT ln K° 2 NO2(g) N2O4(g) • Keq° = 6.75 [products]m[N2O4] [reactants]n[NO2]2 Keq= = ΔG° = -8.314 JK-1 ∙ 298K ∙ ln(6.75) - + ΔG° = -4730 J or 4.73kJ

30. Reactions that Reach Equilibrium ? ΔG° = −RT ln K° PbCl2(s) Pb2+(aq) + 2Cl-(aq) [products] [Pb+2][Cl-]2 [reactants] [PbCl2] Keq= = • Keq° = 1.7 x 10-5 What is the ∆G (positive or negative?) of the above dissolution equation, at 25°C? ∆G = +27.3 kJ