<p>AP Chemistry Chapt. 6 Review 2012-2013</p><p>Objectives:  Describe the properties/characteristics of endothermic and exothermic reactions.  Explain the sign conventions in the equation: ∆E = q+ w, and use the equation to solve problems involving changes in energy, heat and work.  Explain the meaning of a thermochemical equation  Describe how changes to an equation affect the ∆H value of the equation  Use a thermochemical equation to convert between mass and heat  Explain the difference between heat capacity and specific heat; convert between the two  Predict how a substance’s specific heat would influence changes in temperature  Carry out heat calculations using the formula: q = ms∆T  Define calorimetry  Analyze and solve constant-volume (bomb) and constant-pressure calorimetry problems  Define: standard enthalpy of formation  Identify a standard enthalpy of formation reaction  Define: standard enthalpy of reaction  Solve standard enthalpy problems using both the direct method and indirect method (Hess’s law)  Define the following terms: enthalpy of solution, heat of hydration, lattice energy, heat of dilution  Explain the relationship between enthalpy of solution, heat of hydration, and lattice energy; and use this relationship to perform calculations. </p><p>1. What is the change in energy of a system that has 682 J of work done on it and gives off 485 J of heat?</p><p>2. Which is not true for an endothermic reaction?</p><p> a. The temperature of the surroundings decreases. b. The enthalpy change for the reaction is positive. c. Heat flows from the surroundings into the system. d. The products have a lower enthalpy than the reactants. e. All of the above are true.</p><p>3. What are the signs for q and w if a system absorbs 180 J of heat energy while expanding against a constant pressure?</p><p>4. What is the change in internal energy for a system that performs 213 kJ of work on its surroundings and loses 79 kJ of heat in the process?</p><p>5. A 45.0 mL sample of water is heated from 15.0°C to 35.0°C. How many joules of energy have been absorbed by the water? </p><p>6. If 5430 J of energy is used to heat 1.25 L of room temperature water (23.0°C), what is the final temperature of the water?</p><p>7. Given the thermochemical equation: 2SO2 (g) + O2 (g) → 2 SO3 (g) ∆H = -198.2 kJ/mol, calculate the change in enthalpy when 1 mole of SO3 decomposes. o 8. Given: 2 Na2O2(s) + 2 H2O(l) → 4 NaOH(s) + O2(g) ∆H = -126kJ</p><p>Calculate the amount of heat that is evolved (in kJ) from the reaction of 25.0g of Na2O2 with excess water.</p><p>9. If the same amount of heat is added to 5.00 g samples of each of the metals below, which metal will experience the greatest temperature change?</p><p>Metal Specific Heat Capicity (J/g ・ °C) Al 0.897 Au 0.129 Cu 0.385 Fe 0.449 K 0.753</p><p>10. A piece of silver of mass 362 g has a heat capacity of 85.7 J/°C. What is the specific heat of silver?</p><p>11. What is the heat capacity of 60.0 g of water?</p><p>12. A 2.885 g sample of methanol, CH3OH, is combusted in a bomb calorimeter. The temperature of the calorimeter increases by 11.38 °C. If the heat capacity of the bomb is 727.1 J/°C and it contains 1.200 kg of water, what is the heat evolved per mole of methanol combusted? The specific heat of water is 4.184 J/g °C and the molar mass of methanol is 32.04 g/mol.</p><p>13. A 1.96g sample of titanium was burned in a bomb calorimeter that had a heat capacity of 9.84 kJ/°C. The temperature of the calorimeter increased from 36.84°C to 98.82°C. Calculate the heat (kJ) that is released from the combustion of one mole of titanium.</p><p>14. When 50.0 mL of 1.00 M HCl is mixed with 50.0 mL of 1.00 M NaOH (both at 23.0°C), the resulting solution increases in temperature to 29.8°C. Assuming that the solution has the density and specific heat of pure water, calculate the enthalpy of the reaction.</p><p>15. Which of the following chemical equations does not correspond to a standard molar enthalpy of formation? a) Ca(s) + C(s) + 3/2 O2(g) → CaCO3(s) b) C(s) + 1/2 O2(g) → CO(g) c) H2(g) + 1/2 O2(g) → H2O(l) d) N2(g) + 2 O2(g) → N2O4(g) e) SO2(g) + 1/2 O2(g) → SO3(g)</p><p>16. Consider the following thermochemical equations: Fe2O3 + 3CO → 2Fe + 3CO2 ∆H° = -28.0kJ 3Fe +4CO2 → 4CO + Fe3O4 ∆H° = +12.5kJ</p><p>Calculate the value of ∆Ho (in kJ) for 3Fe2O3 + CO → CO2 + 2Fe3O4 17. Iron ore can be converted to iron metal with CO gas. </p><p> o 18. Calculate ∆H rxn for: 4NH3(g) + 5O2(g) → 4NO(g) + 6H2O(l) , given Substance ∆H o f (kJ/mol) H2O(l) -286 NO(g) 90 NH3(g) -46</p><p>-1 19. Knowing that the Hcombusion of octane, C8H18, is -5508.9 kJ·mol calculate the Hf of octane. chemical CO2(g) H2O(g) C8H18(l) -1 -1 -1 Hf -393.5 kJ·mol -241.8 kJ·mol ??? kJ·mol</p><p>20. What volume of CH4(g), measured at 25°C and 745Torr, must be burned in excess oxygen to release 1.00 x 106 kJ of heat to the surroundings? CH4(g) + 2O2(g) → CO2(g) + 2H2O(l) ∆H = -890.3 kJ</p><p>21. If silver nitrate has a lattice energy of -820 kJ/mol and a heat solution of +22.6 kJ/mol, then calculate the heat of hydration for silver nitrate.</p><p>22. Lattice energy of NaI is 686kJ/mol, and enthalpy of hydration is -694 kJ/mol. What is the enthalpy of solution per mol NaI?</p>