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Chapter 17 Thermochemistry 17.1 the Flow of Energy – Heat at Work 17.2 Measuring and Expressing Enthalpy Changes 17.3 He

Chapter 17 Thermochemistry 17.1 the Flow of Energy – Heat at Work 17.2 Measuring and Expressing Enthalpy Changes 17.3 Heat in Changes of State 17.4 Calculating Heats of Reaction. 17.1 The flow of energy – heat at work In what direction does heat flow?

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Chapter 17 Thermochemistry 17.1 the Flow of Energy – Heat at Work 17.2 Measuring and Expressing Enthalpy Changes 17.3 He

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  1. Chapter 17 Thermochemistry 17.1 the Flow of Energy – Heat at Work 17.2 Measuring and Expressing Enthalpy Changes 17.3 Heat in Changes of State 17.4 Calculating Heats of Reaction

  2. 17.1 The flow of energy – heat at work • In what direction does heat flow? • What happens in endothermic and exothermic processes? • In what units is heat flow measured? • On what two factors does the heat capacity of an object depend?

  3. Energy Transformations: • Energy is the capacity for doing work or supplying heat. • Energy has neither mass or volume. • Energy is detected by its effect. • Thermochemistry – is the study of energy changes that occur • during chemical reactions and changes in state. • 5. Chemical Potential Energy – The energy stored in the chemical • bonds of a substance. • 6. Energy changes occur as either heat transfer or work, or a • combination of both. • 7. Heat – symbol = q – energy transfers from one object to another • from higher amounts of energy to less from warm to cold

  4. Exothermic and Endothermic Processes: • Chemical reactions either release or absorb heat. • The “system” is the focus of a study and concern is the flow of heat • either to or from the system. • 3. The “surroundings” is the area around the system that gives or takes • heat from the system. • 4. Law of Conservation of Energy – energy can not be created or • destroyed, the energy however can change form. • 5. Endothermic – The system gains heat, cooler than surrounding, +q • 6. Exothermic – The system loses heat, hotter than surrounding, -q

  5. Units for measuring Heat: • Heat flow is measured two ways. • a. calorie • b. joule • 2. A calorie (cal) is the quantity of heat needed to raise the • temperature of 1g of water 1O C • 3. The word calorie is written with a small “c” • 4. Dietary Calorie is written with a large “C” • 5. 1Calorie = 1000calories • 6. Joule – the SI unit for energy • a. 1J = 0.2390 cal • b. 4.184 J = 1cal

  6. Heat Capacity and Specific Heat: • Heat Capacity – the amount of heat needed to increase the • temperature of an object exactly 1OC • 2. The heat capacity of an object depends on both its mass and • chemical composition. • 3. Specific Heat – is the amount of heat it takes to raise the • temperature of 1g of a substance 1OC • 4. Water has a large specific heat. What does that mean? • 5. To calculate Specific heat • C = q/m x DT • 6. What are the units of specific heat?

  7. In what direction does heat flow? What happens in endothermic and exothermic processes? In what units is heat flow measured? On what two factors does the heat capacity of an object depend?

  8. 17.2 Measuring and Expressing Enthalpy Changes: • What basic concepts apply to calorimetry? • How can you express the enthalpy change for a reaction in a • chemical equation?

  9. Calorimetry: • Is the precise measurement of the heat flow into and out of a system • for chemical or physical processes. • The heat released by the system is equal to the heat absorbed by • its surroundings • Calorimeter – a device used to measure the release or absorption of • heat in chemical or physical processes. • Enthalpy – (H) The heat content of a system at constant pressure • Change in enthalpy – (DH) The heat released or absorbed by a reaction • at constant temperature. q = DH • Equations: • qsurr = m x C x DT • qsys = DH = -qsurr = -m x C x DT • The sign of DH is negative for an exothermic reaction and positive for • an endothermic reaction.

  10. Constant – Volume Calorimeters: • Bomb Calorimeter – a device used for calorimetry experiments in • which the volume remains constant.

  11. Thermochemical Equations: • In a chemical equation, the enthalpy change for the reaction can • be written as either a reactant or a product. • 2. Thermochemical Equation: an equation that includes the • enthalpy change. • 3. Heat of Reaction: is the enthalpy change for the chemical • equation as it is exactly written. • 4. Problems involving enthalpy changes are similar to stoichiometry • problems. • 5. Heat of Combustion: is the heat of reaction for the complete • burning of one mole of a substance.

  12. What basic concepts apply to calorimetry? How can you express the enthalpy change for a reaction in a chemical equation?

  13. 17.3 Heat in Changes of State: • How does the quantity of heat absorbed by a melting solid compare to • the quantity of heat released when the liquid solidifies? • -How does the quantity of heat absorbed by a vaporizing liquid compare • to the quantity of heat released when the vapor condenses? • -What thermochemical changes can occur when a solution forms?

  14. Heats of Fusion and Solidification: • All solids absorb heat as they melt to become liquids. • A gain of heat causes a change in state instead of a change • in temperature. • 3. Molar Heat of Fusion: the heat absorbed by one mole of a solid as • it melts to a liquid at a constant temperature. DHfus • 4. Molar Heat of Solidification: the heat lost when one mole of a • liquid solidifies at a constant temperature. DHsolid • 5. The quantity of heat absorbed by a melting solid is exactly the • same as the quantity of heat released when the liquid solidifies. • D Hfus = - DHsolid

  15. Heats of Vaporization and Condensation: • When liquids absorb heat at their boiling points, they become vapors. • Molar Heat of Vaporization: the amount of heat necessary to vaporize • one mole of a given liquid. DHvap • 3. When vapor condenses, heat is released. • 4. Molar Heat of Condensation: the amount of heat released when 1 mole • of vapor condenses at the normal boiling point. DHcond • 5. The quantity of heat absorbed by a vaporizing liquid is exactly the • same as the quantity of heat released when the vapor condenses. • DHvap = -DHcond

  16. Heat of Solution: • During the formation of a solution, heat is either released or • absorbed. • 2. Good examples; hot and cold packs • 3. Molar Heat of Solution: the enthalpy change caused by the • dissolution of one mole of a substance.

  17. -How does the quantity of heat absorbed by a melting solid compare to the quantity of heat released when the liquid solidifies? -How does the quantity of heat absorbed by a vaporizing liquid compare to the quantity of heat released when the vapor condenses? -What thermochemical changes can occur when a solution forms?

  18. 17.4 Calculating Heats of Reaction: -What are the two ways that you can determine the heat of reaction when it cannot be directly measured?

  19. Hess’s Law: • Hess’s Law: allows you to determine the heat of reaction indirectly • so that we can determine the enthalpy change. • 2. Hess’s Law of Heat Summation: states that if you add two or more • thermochemical equations to give a final equation, then you also • add the heats of reaction to give the final reaction.

  20. Standard Heats of Formation: • Enthalpy changes depend on conditions of the process. • Standard State: a common set of conditions used as a reference • point that refer to the stable form of a substance at 25OC • and 101.3kPa • 3. Standard Heat of Formation: is the change in enthalpy that • accompanies the formation of one mole of a compound from its • elements with all substances in their standard states. DHfO • 4. The standard heat of formation for diatomic molecules = 0 • 5. For a reaction that occurs at standard conditions, you can • calculate the heat of reaction by using standard heats of • formation and is called the standard heat of reaction. • DHO = D HfO (products) - DHfO (reactants)

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