1. Chapter 6

2. Objectives 3.0 Define key terms and concepts. 3.10 Define and apply the First Law of Thermodynamics. 3.11 Calculate the energy produced by a chemical reaction. 3.12 Determine if a reaction is exothermic or endothermic. 3.13 Write and manipulate thermochemical reactions.

3. Energy The capacity to do work Energy = Force x Distance

4. Types of Energy Radiant Energy Energy that comes from the sun Thermal Energy Energy produced by the random motion of atoms and molecules. Measured by temperature Chemical Energy Energy stored in the structure of chemical substances Kinetic Energy The energy of motion Potential Energy Energy available because of an objects position

5. Law of Conservation of Energy Energy can be converted from one form to another, just like matter can be converted Law of Conservation of Energy The total amount of energy in the universe is assumed to be constant.

6. Energy in Chemical Reactions Chemical reactions can either take in or produce energy, typically in the form of heat Thermochemistry The study of heat change in chemical reactions System What we are studying Surroundings Everything that we are not studying

7. Types of Systems Open System Can exchange mass and energy with it’s surroundings

8. Types of Systems Closed System Allows for the transfer of heat, but not mass

9. Types of Systems Isolated System Doesn’t allow for the transfer of heat or mass

10. Types of Thermal Reactions Exothermic Reaction When energy is produced in the reaction Endothermic Reaction When energy is a reactant (taken in)

11. Thermodynamics The study of the transfer of heat energy or the conversion of heat energy and other forms of energy State of a System The values of relevant macroscopic properties such as composition, energy, temperature, and volume. State Functions Properties determined by the state of the system, no matter how it was achieved. The magnitude of change in a state function is only dependent on the initial and final states, not in how the change occurred.

13. The First Law of Thermodynamics Energy cannot be created or destroyed, only converted from one form to another. Because there can be no loss of energy in a system, it can also be said that ΔU = U f – U i ΔU sys + ΔU surr = 0

14. The First Law of Thermodynamics Where q is the sum of the heat exchange between the system and the surrounds q is positive for an endothermic process q is negative for an exothermic process Where w is the work done w is positive for work done on the system w is negative for work done by the surroundings on the system ΔU = q + w

15. Work Typically, work is defined as force x distance; however, in thermodynamics, there are a lot of different types of work taken into account, usually involving work done by a gas on its surroundings. In this formula, the pressure is the equivalence of the force and the distance is the volume. Work is expressed in L x atm, also known as joules (J). Work is not a state function. It is dependent on the process, not just the initial and final states. w =-PΔV

17. Heat The energy that flows into out of a system because of a difference in temperature between a system and its surroundings. For a given system, if q increases, w will decrease and visa versa to maintain ΔU. Is not a state function.

18. The process of dissolving ammonium nitrate in water is an endothermic process. Based on this information, what is the sign for q? Would the flask feel hot or cold? If the reaction of nitrogen gas with oxygen gas to produce nitrogen dioxide produces 66.4kJ of heat, is the process endothermic or exothermic? What is the value of q?

19. If the decomposition of ozone to oxygen gas produces heat, is the reaction exothermic or endothermic? What would the sign for q be? If the production of hydrogen cyanide as indicated below produces 939kJ of heat, what is the value of q? 2NH 3(g) + 3O 2(g) + 2CH 4  2HCN (g) + 6H 2 O (g) Is the reaction endothermic or exothermic?

20. The work done when hydrogen gas is compressed in a canister is 263J. If 52J of heat energy is absorbed by the surroundings during this process, what is the total energy change?

21. If the change of internal energy for a system is 1.5kJ, how many joules of work are conducted if 550J of heat are produced?

22. If a gas absorbs 200J of heat energy and is compressed from 20L to 10L by an opposing pressure of 2atm, what is the internal energy change for this process?

23. What are your questions?

24. Enthalpy of Chemical Reactions Enthalpy (H) Related to the heat of a reaction Used to obtain the amount of heat absorbed or evolved in a chemical reaction Can not be directly determined. Can only calculate relative to arbitrary references. ΔH = ΔU + PΔV

25. Enthalpy of Chemical Reactions Enthalpy of Reaction (ΔH) The difference between the enthalpies of the products and the enthalpies of the reactants Positive for an endothermic reaction Negative for an exothermic reaction

26. Enthalpy of Chemical Reactions The change in enthalpy for a given substance is typically measured per mole of the substance Ice requires 6.01kj of energy per 1 mole to melt Thermochemical Reactions Shows the molar relationships in a chemical reaction as well as the change in enthalpy 2Na (s) + 2H 2 O (l)  2NaOH (aq) + H 2(g) ΔH = -367.5kJ

27. Writing Thermochemical Equations Always specify the physical states of matter since they help determine the enthalpy change. If a chemical reaction is multiplied by a factor of n, then ΔH must be multiplied by that same number. If we reverse the direction of the chemical reaction, then the sign for ΔH must be changed.

28. Aqueous sodium hydrogen carbonate solution reacts with hydrochloric acid to produce aqueous sodium chloride, water and carbon dioxide gas. The reaction absorbs 11.8kJ of heat at a constant pressure for each mole of sodium hydrogen carbonate. What is the thermochemical reaction?

29. When ethanol (C 2 H 5 OH (l) ) reacts with oxygen gas, carbon dioxide and water are produced. The reaction produces 1.3kJ of heat at a constant pressure for each mole of ethanol reacted. What is the thermochemical reaction? What would be the total heat produced if 3 moles of ethanol were reacted? What is the reverse of this reaction?

30. When 2 moles of hydrogen gas and 1 mole of oxygen gas reacts to produce liquid water, 572kJ of heat is produced. Write the balanced thermochemical reaction. How much heat would be produced if only 1 mole of water is produced? Write the reverse reaction.

31. Stoichiometry and Thermochemistry The amount of heat that is produced in a reaction is dependent on the amount of reactant. Grams of Compound Moles of Compound Amount of Heat Generated

32. How much heat is evolved when 831kg of ammonia is produced from the reaction of nitrogen and hydrogen gas as illustrated below: N2 (g) + 3H 2(g)  2NH 3(g) ΔH = -91.8kJ

33. How much heat is evolved when 89.6kg of hydrogen gas is reacted with nitrogen gas to produce ammonia as illustrated below: N 2(g) + 3H 2(g)  2NH 3(g) ΔH = -91.8kJ

34. Nitric oxide combines with oxygen to produce nitrogen dioxide. 2NO (g) + O 2(g)  2NO 2(g) ΔH = -114kJ How much heat evolves if 159.4g of NO gas is reacted with excess oxygen gas?

35. Nitric oxide combines with oxygen to produce nitrogen dioxide. 2NO (g) + O 2(g)  2NO 2(g) ΔH = -114kJ What is the enthalpy change per gram of NO?

36. Methane reacts with oxygen gas to produce carbon dioxide and water as shown in the following reaction CH 4(g) + 2O 2(g)  CO2 (g) + 2H 2 O (l) ΔH = -890.3kJ How much energy is produced if 12.4g of methane is reacted with 21.9g of oxygen gas?

37. Comparison of ΔH and ΔU The internal energy change (ΔU) of a gaseous system can also be calculated assuming ideal gas behavior and constant temperature. Where R is 8.314J/K·mole Where Δn = # moles of product gases - # moles of reactant gases ΔU = ΔH - RTΔn

38. What is the change in internal energy when 2 moles of hydrogen gas are reacted with oxygen gas to produce 2 moles of water at 2.1atm and 20°C. 2H 2(g) + O 2(g)  2H 2 O (l) ΔH = -572kJ

39. What is the change in internal energy when 1 mole of nitrogen gas are reacted with 3 moles of hydrogen gas to produce 2 moles of ammonium at 1.4atm and 13°C. N2 (g) + 3H 2(g)  2NH 3(g) ΔH = -91.8kJ

40. Using the reaction for the formation of ammonia gas provided below, what is the change in internal energy if this reaction is conducted at 1.8atm and 21°C? N2 (g) + 3H 2(g)  2NH 3(g) ΔH = -91.8kJ

41. What are your questions?

42. Standard Enthalpy of Formation Standard Enthalpy of Formation (ΔH° f ) The heat of change that results when 1 mole of the compound is formed from its elements at a pressure of 1atm. The reference point from which enthalpy is measured The most stable form of each of the elements (O 2, Cl 2, Al, graphite, etc)

44. Standard Enthalpy of Reaction Standard Enthalpy of Reaction (ΔH° rxn ) The enthalpy of a reaction carried out at 1atm. Direct Method Indirect Method Hess’s Law The change in enthalpy for a reaction is the same no matter if the reaction occurs in one step or many steps Δ H° rxn = ΣΔH° f(products) - ΣΔH° f(reactants)

45. What is the standard enthalpy change for the reaction of ammonia with oxygen gas as shown below? NH 3(g) + O 2(g)  NO (g) + H 2 O (g)

46. Calculate the enthalpy change for the following reaction using standard enthalpies of formation. NO 2(g) + H 2 O (l)  HNO 3(aq) + NO (g)

47. Hydrazine, a component of rocket fuel, is formed from the reaction of nitrogen gas and hydrogen gas in the reaction shown below: N 2(g) + H 2(g)  N 2 H 4(g) Using the following thermochemical reactions, calculate the enthalpy change for the formation of hydrazine. N 2 H 4(g) + O 2(g)  N 2(g) + 2H 2 O (l) ΔH = -622.2kJ H 2(g) + ½O 2(g)  H 2 O (l) ΔH = -285.8kJ

48. Ammonia will burn in the presence of a platinum catalyst to produce nitric oxide. NH 3(g) + O 2(g)  NO (g) + H 2 O (g) Using the following thermochemical reactions, calculate the heat of reaction at a constant pressure. N 2(g) + O 2(g)  2NO (g) ΔH = 180.6kJ N 2(g) + 3H 2(g)  2NH 3(g) ΔH = -91.8kJ 2H 2(g) + O 2(g)  2H 2 O (g) ΔH = -483.7kJ

49. Calculate the standard enthalpy change, ΔH o, for the formation of 1 mol of strontium carbonate (the material that gives the red color in fireworks) from its elements.

50. What are your questions?

51. Calorimetry The measure of heat changes Measured using a closed system called a calorimeter

52. Calorimetry Specific Heat (s) Amount of heat required to raise the temperature of one gram of a substance 1°C.

53. Calorimetry Heat Capacity (C) The amount of heat required to raise the temperature of a given quantity of a substance 1°C. Where m is the mass of the compound in grams Where s is the specific heat of the substance. C = ms

54. Calorimetry The amount of heat that was absorbed or released during a specific process using the following equations WhereΔt = t final – t initial q is positive for an endothermic process q is negative for an exothermic process q = msΔt q = CΔt

55. Constant-Pressure Calorimetry Used to determine heat change for non- combustion reactions Is an isolate system Since the pressure is constant, q rxn = ΔH

56. A 200g sample of water is heated from 2.25°C to 29.1°C. How much heat was absorbed in kJ by the water?

57. If a 153g sample of water is cooled from 9.9°C to 1.5°C. How much heat was lost in kJ by the water?

58. A 8.7g sample of a metal requires 675J of heat to raise its temperature from 20°C to 29°C. What is the heat capacity of the metal?

59. Calculate the specific heat of a liquid that requires 428J to raise a 15.7g sample of the liquid by 12°C.

60. 500mL of 0.5M sulfuric acid is reacted with 750mL of 0.65M acetic acid in a constant pressure calorimeter. The initial temperature of both solutions is 21.5°C. The temperature of the solution after the reaction is complete is 23.1°C. Calculate the molar heat of change for each of the chemicals in this reaction? Assume the densities and specific heat of both reactants is that of water.

61. A constant volume calorimeter contains 150mL of water with an initial temperature of 22.5°C. A 15.6g piece of metal with an initial temperature of 90.1°C is placed in the water in the calorimeter. The final temperature of the solution is 51.7°C. Calculate the specific heat of the metal. The specific heat of water is 4.184J/g°C.

62. Constant-Volume Calorimetry Bomb Calorimeter Measures Heat of Combustion Isolate system, so no heat lost to surroundings Where q sys = 0 q cal = C cal Δt q sys = q cal + q rxn

63. If 0.562g of graphite is reacted in a calorimeter with excess oxygen at 25°C and 1atm, the temperature of the water raises to 25.89°C. What is the heat of the reaction if the heat capacity of the calorimeter and its contents is 20.7kJ/°C?

64. A 2.84g sample of ethanol is burned in the presence of excess oxygen in a bomb calorimeter and the temperature of the water rises from 25°C to 33.73°C. If the heat capacity of the calorimeter and contents is 9.63kJ/°C, what is the enthalpy change for the reaction of 1 mole of ethanol? C 2 H 5 OH (l) + 2O 2(g)  2CO 2(g) + 3H2O (l)

65. What are your questions?