Describe the Properties/Characteristics of Endothermic and Exothermic Reactions

AP Chemistry Chapt. 6 Review 2012-2013

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.

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?

2. Which is not true for an endothermic reaction?

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.

3. What are the signs for q and w if a system absorbs 180 J of heat energy while expanding against a constant pressure?

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?

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?

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?

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

Calculate the amount of heat that is evolved (in kJ) from the reaction of 25.0g of Na2O2 with excess water.

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?

Metal Specific Heat Capicity (J/g ・ °C) Al 0.897 Au 0.129 Cu 0.385 Fe 0.449 K 0.753

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?

11. What is the heat capacity of 60.0 g of water?

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.

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.

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.

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)

16. Consider the following thermochemical equations: Fe2O3 + 3CO → 2Fe + 3CO2 ∆H° = -28.0kJ 3Fe +4CO2 → 4CO + Fe3O4 ∆H° = +12.5kJ

Calculate the value of ∆Ho (in kJ) for 3Fe2O3 + CO → CO2 + 2Fe3O4 17. Iron ore can be converted to iron metal with CO gas.

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

-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

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

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.

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?