New Material Exam In any chemical process, energy must be conserved. This is a statement of the______________ law of thermodynamics. At what temperatures will a reaction be spontaneous if ΔrH° = +60.9 kJ and ΔrS° = +294 J/K? A. Temperatures between 179 K and 235 K. B. The reaction will be spontaneous at any temperature. C. All temperatures below 207 K. D. All temperatures above 207 K. E. The reaction will never be spontaneous. Which one of the following processes involves an increase in the entropy of the system? A. CaO(s) + CO2(g) → CaCO3(s) B. 3 H2(g) + N2(g) → 2 NH3(g) C. 2 Na(s) + Cl2(g) → 2 NaCl(s) D. N2(g) + 2 O2(g) → 2 NO2(g) E. 2 Ag2O(s) → 4 Ag(s) + O2(g) Thermodynamics can be used to determine all of the following EXCEPT A. .the temperature at which a reaction is spontaneous. B. the entropy change of a reaction. C. the extent to which a reaction occurs. D. the rate of reaction. E. the direction in which a reaction is spontaneous. The standard free energy change for a chemical reaction is +13.3 kJ/mol. What is the equilibrium constant for the reaction at 117 °C? (R = 8.314 J/K×mol) A. 0.996 B. 1.15E-6 C. 1.66E-2 D. 60.4 E. 7.94E-5 Calculate ΔrG° for the reaction below at 25.0 °C CS2(g) + 3 Cl2(g) → S2Cl2(g) + CCl4(g) given ΔrH° = –231.1 kJ/mol-rxn and ΔrS° = –287.6 J/K×mol-rxn. A. –145.3 kJ/mol-rxn B. +56.5 kJ/mol-rxn C. –316.9 kJ/mol-rxn D. –518.7 kJ/mol-rxn E. –56.5 kJ/mol-rxn The following anions can be separated by precipitation as silver salts: Cl–, Br–, I–, CrO42–. If Ag+ is added to a solution containing the four anions, each at a concentration of 0.10 M, in what order will they precipitate? Compound Ksp A. AgCl → Ag2CrO4 → AgBr → AgI AgCl 1.8E-10 B. AgI → AgBr → AgCl → Ag2CrO4 Ag_2_CrO4 1.1E-12 C. Ag2CrO4 → AgI → AgBr → AgCl AgBr 5.4E-13 D. AgI → AgBr → Ag2CrO4 → AgCl AgI 8.5E-17 E. Ag2CrO4 → AgCl → AgBr → AgINew Material Exam If a chemical reaction occurs in a direction that has a positive change in entropy then A. the reaction must be exothermic. B. the change in enthalpy must be negative. C. heat goes from the system into the surroundings. D. the reaction must be spontaneous. E. the disorder of the system increases. Assuming the following reaction proceeds in the forward direction, 3 Sn4+(aq) + 2 Cr(s) → 3 Sn2+(aq) + 2 Cr3+(aq) A. Sn4+(aq) is the reducing agent and Sn2+(aq) is the oxidizing agent. B. Sn4+(aq) is the reducing agent and Cr(s) is the oxidizing agent. C. Cr(s) is the reducing agent and Cr3+(aq) is the oxidizing agent. D. Cr(s) is the reducing agent and Sn4+(aq) is the oxidizing agent. E. Cr(s) is the reducing agent and Sn2+(aq) is the oxidizing agent. Write a balanced half-reaction for the reduction of CrO42–(aq) to Cr(OH)3(s) in a basic solution. A. CrO42–(aq) + 4 H2O( ) + 3 e– → Cr(OH)3(s) + 5 OH–(aq) B. CrO42–(aq) + 3 H+(aq) + 3 e– → Cr(OH)3(s) C. CrO42–(aq) + 3 OH–(aq) → Cr(OH)3(s) + 2 O2(g) D. CrO42–(aq) + 3 OH–(aq) + 3 e– → Cr(OH)3(s) + 2 O2(g) E. CrO42–(aq) + 3 H+(aq) → Cr(OH)3(s) + 2 e– The concentration of Pb2+ in an aqueous solution is 2.49E-3 M. What concentration of SO42– is required to begin precipitating PbSO4? The Ksp of PbSO4 is 2.5E-8 A. 1.20E-4 M B. 1.00E-5 M C. 8.37E-7 M D. 9.96E+4 M E. 3.17E-3 M Write a balanced chemical equation for the overall reaction represented by the cell notation below. Zn(s) | Zn2+(aq) || H+(aq) | H2(g) | Pt(s) A. H2(g) + Zn(s) → 2 H+(aq) + Zn2+(aq) B. 2H+(aq) + Zn(s) → ZnH2(s) C. 2H+(aq) + Zn(s) → H2(g) + Zn2+(aq) D. H2(g) + Zn2+(aq) → 2 H+(aq) + Zn(s) E. 2H+(aq) + Zn2+(aq) → H2(g) + Zn(s) Use the following thermodynamic data. Species ΔfH° (kJ/mol) S° (J/K×mol) Fe(s) 0.0 27.8 O2(g) 0.0 205.1 Fe3O4(s) –1118.4 146.4New Material Exam to calculate ΔS° (universe) for the formation of Fe3O4(s) at 298.15 K. 3 Fe(s) + 2 O2(g) → Fe3O4(s) A. +3404 J/K B. –3404 J/K C. +561.2 J/K D. +7639 J/K E. –1162 J/K Write a balanced half-reaction for the reduction of NO3–(aq) to NO(g) in an acidic solution. A. NO3–(aq) + 4 H+(aq) + 3 e– → NO(g) + 2 H2O( ) B. NO3–(aq) + 4 H+(aq) → NO(g) + 2 H2O( ) C. NO3–(aq) + 4 e– → NO(g) + O2(g) D. NO3–(aq) + 4 H+(aq) → NO(g) + 2 H2O( ) + 2 e E. NO3–(aq) + 3 e– → NO(g) +