Survey
* Your assessment is very important for improving the workof artificial intelligence, which forms the content of this project
* Your assessment is very important for improving the workof artificial intelligence, which forms the content of this project
CHEM102 December 11, 2003 Review for the Final Exam Structure of the final exam. 6 problems + 1 extra credit problem 1. 2. 3. 4. 5. 6. 7. Short answers. Definitions and concepts Thermodynamic calculations Nomenclature of organic compounds Functional groups in organic compounds; hybridization Chemical equilibria: calculations of equilibrium constants; LeChatelier’s principle Combined question: rates of chemical reactions, reaction mechanisms and thermodynamics Extra credit problem: generic 2. Example of Problem 2. (see problem set on Thermodynamics from December 10) 1. Consider the following chemical reaction: N2(g) + 2 O2(g) → 2NO2(g) a) Using the tables provided below, calculate the standard enthalpy, standard entropy and standard free energy change (∆ ∆Ho, ∆So, and ∆Go, respectively) for the above reaction at 60 °C: ∆H0 = 2 ∆H0f(NO2(g))-( ∆H0f(N2(g)) + 2∆H0f(O2(g)) = 2 x 33.84-2 x 0 – 2 x 0 = 67.68 kJ/mol ∆S0 = 2 S0(NO2(g))-( S0(N2(g)) + 2S0(O2(g)) = 2 x 240.45 – 191.50-2 x 205.0 =-120.5 J/mol-K ∆G0 = ∆H0 - T ∆S0 = 67.68 - 333 (-0.1205) = 107.807 kJ/mol b) Calculate the equilibrium constant for this reaction at 60 oC. ∆G0 = -RT lnK; K = e-∆G0/RT= e-107807/8.31x333 = 1.2 x 10-17 very small number- equilibrium favors reactants at this temperature c) Will an increase in temperature produce an increase or decrease in the mole fraction of NO2(g) at equilibrium? Explain. The reaction is endothermic (∆H0<0); increasing temperature will favor formation of products d) At what temperature if any will this reaction have an equilibrium constant equal to 1? You may assume that ∆H°° and ∆S°° are temperature-independent. K = 1 occurs when ∆G° = 0; therefore, ∆H° - T∆S° = 0, or T = ∆H° / ∆S° = 67680/(-120.5) = -561.66 K , which physically means that there is no temperature at which K can be equal to 1. In other words, both the enthalpic and entropic terms are unfavorable (endothermic reaction with negative entropy change- always non spontaneous at standard conditions). 3. Example of problem 3. (see workshop 2) Give the IUPAC names for each of the following compounds: 1 CHEM102 December 11, 2003 CH3 CH3CH2 CH2CH3 HC C CH2Cl CH3 CH3CH2 CCH2CCH3 CH2CHCH2CH3 C a) CH3 4,4-Dimethyl-1-hexyne H c) b) C H 1-Chloro-2-methyl-2-phenyl-butane or (1-chloro-2-methyl)-2-butylbenzene Cis-6-methyl-3-octene 4. Example of problem 4. The compound whose structure is shown below is acetylsalicylic acid, better known as aspirin: O OH O O CH3 a) Identify the functional groups in aspirine: carboxyl O OH phenyl O O CH3 ester b) Give the hybridization of each carbon atom in the aspirin molecule: O OH O O CH3 2 all carbons except for the CH3 carbons- sp hybridized; CH3 carbons- sp3 hybridized 2 CHEM102 December 11, 2003 5. Example of problem 5. Consider the following reaction at equilibrium: A(g) V 2B (g) The equilibrium constant of the reaction was measured at several temperatures and the following data were obtained: Temperature (ºC) [A], M [B], M 200 0.0125 0.843 300 0.171 0.764 400 0.250 0.724 a) Based on the data, calculate the equilibrium constant (Kc) at each temperature. Kc = [B]2/[A] At 200 oC, Kc = 0.8432/0.0125 = 56.85 At 300 oC, Kc = 0.7642/0.171 = 3.41 At 400 oC, Kc = 0.7242/0.250 = 2.10 b) Is the reaction exothermic or endothermic? Explain your answer Since the equilibrium constant decreases with increasing the temperature, the reaction is exothermic, according to the LeChatelier’s principle (formation of reactants is favored at higher temperatures). 6. Example of problem 6. Consider the following reaction: H2(g) + 2 ICl (g) 2 HCl (g) + I2 (g) The rate law for this reaction is first order in both H2 and ICl (meaning, Rate = k x [H2] x [ICl]). a) Which of the following mechanisms are consistent with the observed rate law? (a) 2ICl(g) + H2(g) 2HCl (g) + I2(g) (termolecular reaction (b) H2(g) + ICl(g) HI(g) + HCl(g) HI(g) + ICl(g) HCl(g) + I2(g) (slow) (fast) (c) H2(g) + ICl(g) HI(g) + HCl(g) HI(g) + ICl(g) HCl(g) + I2(g) (fast) (slow) (d) H2(g) + ICl(g) HICl(g) + H (g) H(g) + ICl(g) HCl(g) + I(g) HICl(g) HCl(g) + I(g) I(g) + I(g) I2(g) (slow) (fast) (fast) (fast) Answer: b) and d) are consistent, since the stoichiometry of the slow (rate-determining) step will yield the rate law equation. 3 CHEM102 December 11, 2003 b) Based on the mechanism that is consistent with the rate law and on the thermodynamic tables provided, calculate the thermodynamic parameters (∆ ∆Hº, ∆Sº, and ∆Gº) of the individual steps, and of the overall reaction at 25 ºC. Is the reaction spontaneous at 25 ºC? For the overall reaction: ∆Hº= 2∆Hºf(HCl (g)) + ∆Hºf(I2 (g)) -∆Hºf(H2 (g)) -2∆Hºf(ICl (g)) = 2 (-92.3) + 62.25 –0 – 2 17.782 = -157.91 kJ/mol ∆Sº= 2Sº(HCl (g)) + Sº(I2 (g)) - Sº(H2 (g)) -2 Sº(ICl (g)) = 2 186.69 + 260.57 –130.58 – 2 247.44 = 8.49 J/mol-K ∆Gº= ∆Hº-T ∆Sº =157910 J/mol – 298.15 K x 8.49 J/mol-K = 155379 J/mol = 155.379 kJ/mol The reaction is not spontaneous in the forward direction at 25 oC (∆Gº>0). c) Calculate the equilibrium constant of the overall reaction at 25 ºC. Which direction will the equilibrium be shifted upon increasing the temperature? Explain your answer. ∆Gº= -RT lnK; K = e-∆Gº/RT= e-155379/(8.31 x 298.15) = 5.8 x 10-28 This is a very small equilibrium constant. Since the overall reaction is exothermic, increasing temperature will favor formation of the reactants, based on the LeChatelier’s principle. Appendix 1. Thermodynamic quantities for selected substances at 298.15 K (25 oC) Substance Carbon C(s, graphite) CO(g) CH3OH(l) CH3COOH(l) Hydrogen H(g) H2(g) Iodine I2(g) I(g) HI(g) ICl(g) Chlorine Cl2(g) HCl(g) Nitrogen N2(g) N(g) NO(g) N2O(g) Oxygen O2(g) O(g) H2O(l) H2O(g) ∆Hof (kJ/mol) ∆Gof (kJ/mol) So (J/mol-K) 0 -110.5 -238.6 -487.0 0 -137.2 -166.23 -392.4 5.69 197.9 126.8 159.8 217.94 0 203.26 0 114.60 130.58 62.25 106.6 25.94 17.782 19.37 70.16 1.30 -5.44 260.57 180.66 206.3 247.44 0 -92.30 0 -95.27 222.96 186.69 0 472.7 90.37 81.6 0 455.5 86.71 103.59 191.50 153.3 210.62 220.0 0 247.5 -285.83 -241.82 0 230.1 -237.13 -228.57 205.0 161.0 69.91 188.83 4 CHEM102 December 11, 2003 Definitions. Define each of the following: 1. 2. 3. 4. 5. 6. 7. 8. 9. 10. 11. 12. 13. 14. 15. 16. 17. 18. 19. 20. 21. 22. 23. 24. 25. 26. 27. 28. 29. 30. 31. 32. 33. Solution, solvent, solute, solubility Saturated solution Molarity Molality Mass percent Mole fraction Henry’s law Raoult’s law Colligative properties. Vapor pressure lowering, boiling-point elevation, freezing-point depression Osmosis, Osmotic Pressure Hypo-, hyper-, and isotonic solutions Organic compounds Hybridization: sp3, sp2, sp σ-bond, π-bond Hydrocarbons Alkanes Alkenes Alkynes Straight-chain hydrocarbons Branched-chain hydrocarbons Structural isomers Cis-trans isomers Functional groups Aromatic compounds Alcohols Ethers Esters Amines Carboxylic Acids Carbonyl group Addition reactions Substitution reactions Markovnikov’s rule 34. 35. 36. 37. 38. 39. 40. 41. 42. 43. 44. 45. 46. 47. 48. 49. 50. Chemical kinetics Reaction rate Rate law Rate constant First order reaction Expression for the first order reaction rate Second order reaction Expression for the second order reaction rate Zero-order reaction Expression for the zero-order reaction rate Half-life of a reaction Expression for half-life of the first-order reaction Expression for half-life of the second-order reaction Collision theory Activation energy Activated complex Arrhenius equation 5 CHEM102 December 11, 2003 51. 52. 53. 54. 55. 56. 57. 58. 59. 60. 61. 62. 63. 64. 65. 66. 67. 68. 69. 70. 71. 72. 73. 74. 75. 76. 77. 78. Reaction mechanism Elementary step Reactants Products Intermediates Catalysts Molecularity of a reaction Unimolecular reaction Bimolecular reaction Termolecular reaction Rate-determining (rate-limiting) step Homogeneous catalysis Heterogeneous catalysis Chemical equilibrium Equilibrium constant: definition, expression for a given chemical reaction Homogeneous equilibrium Heterogeneous equilibrium Reaction quotient Equilibrium position Le Châtelier’s principle Bronsted acid, Bronsted base Conjugate acid-base pairs Strong acid, strong base Weak acid, weak base Acid ionization constant, base ionization constant Ion product of water pH, pOH Neutral, acidic, basic solutions 1. 2. 3. 4. 5. 6. 7. 8. 9. 10. 11. 12. 13. 14. 15. 16. 17. 18. 19. 20. 21. 22. 23. Buffer solution Henderson-Hasselbach equation Solubility product Molar solubility Soubility (in grams per liter) First Law of Thermodynamics Heat of reaction, internal energy, enthalpy Hess’s law State function System, surroundings, Universe Thermochemical equations Standard state Standard enthalpy of formation Standard enthalpy of reaction Hess’s law Second Law of Thermodynamics Entropy Spontaneous processes Entropy of the system, surroundings, Universe Standard entropy of reaction Third Law of Thermodynamics Gibbs Free Energy Equilibrium constant in terms of standard Gibbs Free Energy 6