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UNIT 10

UNIT 10. Thermodynamics – chapter 17 Organic Chemistry –chapters 22 & 24. Thermodynamics. Chapter 17 -. Thermochemistry – Heat and Chemical Change. Observing Heat Flow.

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UNIT 10

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  1. UNIT 10 Thermodynamics – chapter 17 Organic Chemistry –chapters 22 & 24

  2. Thermodynamics Chapter 17 - Thermochemistry – Heat and Chemical Change

  3. Observing Heat Flow • Hook index fingers through each end of rubber band. Without stretching, place it against your upper lip or forehead. Note the temperature of the rubber band. • Move rubber band away from your skin, quickly stretch and hold it, and then place it back against your skin. Note any temp. change. • Fully stretch the band, and then allow it to return to its original shape. Place it against your skin and note any temp. changes.

  4. Think About It • Did the rubber band feel cooler or warmer after it was stretched in Step 2? • Did the rubber band feel cooler or warmer after it returned to its original shape in Step 3? • Think about temperature changes you observed and form some initial answers to the following questions. What is heat? In what direction does heat flow?

  5. THERMOCHEMISTRY study of heat changes that occur during chemical reactions and changes of state

  6. Some important concepts… • Energy = capacity to do work. • Temperature = measure of Kinetic Energy • Chemical potential energy = Energy stored in chemical bonds of a substance.

  7. HEAT • Energy that transfers from one object to another because of a temperature difference. • Always flows from warmer object to a cooler object. • if the objects remain in contact, they will become equal in temperature • represented by “q”

  8. EXO vs. ENDO Exothermic = • the system loses heat as the surroundings heat up (heat is given off) • “q” is negative (-) Endothermic = • the system gains heat as the surroundings cool down (heat is absorbed) • “q” is positive (+)

  9. How to measure Heat Flow • Heat flow is measured in two common units (the calorie and the joule). • calorie (cal) = the amount of heat needed to raise the temperature of 1 g of pure water 1oC. • Important Relationships: • 1 food Calorie = 1000 calories • 4.184 J = 1 cal

  10. Calorimetry The accurate and precise measurement of heat change for chemical and physical processes. The insulated device used to measure the absorption or release of heat in chemical or physical processes is called a calorimeter.

  11. What happens to the HEAT in lava once it flows out of a volcano? • Heat is released • Why does lava cool more quickly in water than on land? • Water has greater capacity to absorb heat than does air; temp. difference between the lava and water may be greater than between lava and air

  12. Heat Capacity and Specific Heat • Heat Capacity • Amount of heat needed to increase the temperature of an object exactly 1oC. • Depends on both the mass and chemical composition. • Specific Heat • Amount of heat needed to raise the temperature of 1 g of the substance 1oC.

  13. What types of things have High Specific Heat? • Water (1 calorie raises 1g 1oC) • What types of things have low specific heat? • Metals (1 calorie raises 1g Fe 9oC) • What is the relationship between specific heat and heat capacity?

  14. Specific Heat Formula Q = m DT Cp • Q = Heat (J or cal) • m = mass (grams) DT = change in Temp. (Tf – Ti) • Cp = Specific Heat (J / goC) or (cal / goC)

  15. Sample Problem Find the specific heat of copper when its weight is 95.4 grams. The temperature went from 50oC to 27oC and the heat capacity is 849 J. 849 J= .387 J/goC 95.4g x 23oC q = 849 J (heat) m = 95.4 (mass) T = 50oC - 27oC = 23oC

  16. Heat Gained • How much heat is required when 49.0 grams of water is heated from 4oC to 43oC? Calculate in Joules. Specific Heat of water = 1 cal/g x oC or 4.18 J/g x oC mass = 49.0 g T = 43oC – 4 oC = 39oC Heat Gained = Specific Heat x m x T Calories = 1.00 cal x 49.0 g x 39oC = 1911 cal g oC

  17. Heat Gained (Problem continued) Joules = 4.18 J x 49g x g oC 39 oC = 7988 J

  18. More Practice • When 435 J of heat is added to 3.4 g of olive oil at 21oC, the temperature increases to 85oC. What is the specific heat of the oil? • How much heat is required to raise the temperature of 250.0 g of mercury 52oC? The specific heat of mercury is 0.14 J/goC. 1.99 J/goC 1820 J

  19. Thermochemical Equation • An equation that includes the amount of heat gained or released in the chemical equation

  20. CaCO3 + 176 kJ CaO + CO2 Endothermic Equations: (Heat used) Exothermic Equations: C + O2 CO2 + 393.5 kJ (Heat released)

  21. Calculate the energy required to decompose 5.20 moles CaCO3. 5.20 moles CaCO3 x 176 kJ = 915 kJ 1 mole

  22. Sample Problem How much heat is released when 3.84 g of SO3 reacts with H2O? SO3 + H2O H2SO4 + 129.6 KJ 3.84 g SO3 x 1 mole x 129.6 KJ = 6.22 KJ 80 g SO3 1 mole

  23. HEAT IN CHANGES OF STATE • Copy picture of Heating Curve for Water • In which region(s) of the graph is temperature constant? • Describe how the amount of energy required to melt a given mass of ice compares to the energy required to vaporize the same mass of water? Explain. • Which region of the graph represents the coexistence of solid and liquid? Liquid and Vapor?

  24. When a material condenses or freezes is energy released or absorbed? Explain your reasoning. • When a material evaporates or melts, is energy released or absorbed? Explain your reasoning.

  25. Energy (as heat) during Changes of State • Compounds absorb energy (heat) when breaking bonds. • Solids absorb energy to melt into liquids. (ice cubes melting) • Liquids absorb energy to evaporate into gases. • Energy goes to changing state (breaking bonds), there is no temperature change!

  26. Compounds release energy when forming bonds. • Gases release energy to condense into a liquid. • Liquids release energy to freeze into a solid. • Energy goes to changing state (making bonds), there is no temperature change!

  27. Equations for Heating Curve • No Phase Change: Q = m(▲T)Cp • Latent Heat of Fusion: (Solid Liquid) Q = m ▲Hfus • Latent Heat of Vaporization: (Liquid Gas) Q= m ▲Hvap

  28. Practice Problem • How much heat (in kJ) is absorbed when 24.8 g water at 100oC is converted to steam at 100oC? • 56,048 J  56.05 kJ • How much heat is released when 10g of water at 0oC is converted to 10g of ice at 0oC?

  29. How much heat (in kJ) is absorbed when 5 g of water at 50oC is converted to steam and heated to 110oC? • How much heat (in KJ) is released when 10g of steam at 115oC is converted to 10g of ice at -5oC ?

  30. molar heat of fusion = heat absorbed by one mole of a substance in melting from a solid to a liquid at constant temperature. • molar heat of solidification = heat lost when one mole of a liquid solidifies at a constant temperature Hfus = - Hsolid

  31. Enthalpy Amount of heat that a substance has at a given temperature and pressure = H H = Change in enthalpy Heat of reaction - heat that is released or absorbed during a chemical reaction H = negative for exothermic(-) H = positive for endothermic(+)

  32. Exothermic Endothermic H is negative H is positive E E ReactantsProducts ReactantsProducts

  33. Hess's Law If you add two or more thermochemical equations to give a final equation, then you can also add the heats of reaction to give the final heat of reaction.

  34. Enthalpy changes occur when a compound is formed from its elements • Diatomic molecules = 0 • Use table 11-6 pg 316 • H is the difference between standard enthalpies of formation of all reactants and products

  35. Formula for Enthalpy Problems H = H f (products) - H f ( reactants) = change H = enthalpy = standard conditions (25oC and 101.3 kpa) f = formation of the compound

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