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Thermochemistry

Thermochemistry

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Thermochemistry

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  1. Thermochemistry David P. White University of North Carolina, Wilmington Chapter 5 Chapter 5

  2. The Nature of Energy Kinetic and Potential Energy From Physics: • Force is a push or pull on an object. • Work is the product of force applied to an object over a distance: w = Fd • Energy is the work done to move an object against a force. • Kinetic energy is the energy of motion: Chapter 5

  3. The Nature of Energy Kinetic and Potential Energy • Potential energy is the energy an object possesses by virtue of its position. • Potential energy can be converted into kinetic energy. Example: a ball of clay dropping off a building. Chapter 5

  4. The Nature of Energy Energy Units SI Unit for energy is the joule, J: We sometimes use the calorie instead of the joule: 1 cal = 4.184 J (exactly) A nutritional Calorie: 1 Cal = 1000 cal = 1 kcal Chapter 5

  5. The Nature of Energy Systems and Surroundings System: part of the universe we are interested in. Surroundings: the rest of the universe. Chapter 5

  6. First Law of Thermodynamics Internal Energy • Internal Energy: total energy of a system. • Cannot measure absolute internal energy. • Change in internal energy, DE = Efinal - Einitial Chapter 5

  7. First Law of Thermodynamics • Relating DE to Heat and Work • Energy cannot be created or destroyed. • Energy of (system + surroundings) is constant. • Any energy transferred from a system must be transferred to the surroundings (and vice versa). • From the first law of thermodynamics: • when a system undergoes a physical or chemical change, the change in internal energy is given by the heat added to or absorbed by the system plus the work done on or by the system: • DE = q + w Chapter 5

  8. First Law of Thermodynamics Relating DE to Heat and Work Chapter 5

  9. First Law of Thermodynamics Relating DE to Heat and Work Chapter 5

  10. First Law of Thermodynamics Endothermic and Exothermic Processes Endothermic: absorbs heat from the surroundings. Exothermic: transfers heat to the surroundings. An endothermic reaction feels cold. An exothermic reaction feels hot. Chapter 5

  11. First Law of Thermodynamics State Functions State function: depends only on the initial and final states of system, not on how the internal energy is used. Chapter 5

  12. First Law of Thermodynamics State Functions Chapter 5

  13. Enthalpy Enthalpy, H: Heat transferred between the system and surroundings carried out under constant pressure. Can only measure the change in enthalpy: DH = Hfinal - Hinitial = qP Chapter 5

  14. Enthalpies of Reaction For a reaction DHrxn = H(products) - H (reactants) Enthalpy is an extensive property (magnitude DH is directly proportional to amount): CH4(g) + 2O2(g)  CO2(g) + 2H2O(g) DH = -802 kJ 2CH4(g) + 4O2(g) 2CO2(g) + 4H2O(g) DH = -1604 kJ When we reverse a reaction, we change the sign of DH: CO2(g) + 2H2O(g)  CH4(g) + 2O2(g) DH = +802 kJ Change in enthalpy depends on state: H2O(g)  H2O(l) DH = -88 kJ Chapter 5

  15. Calorimetry Heat Capacity and Specific Heat Calorimetry = measurement of heat flow. Calorimeter = apparatus that measures heat flow. Heat capacity = the amount of energy required to raise the temperature of an object (by one degree). Molar heat capacity = heat capacity of 1 mol of a substance. Specific heat = specific heat capacity = heat capacity of 1 g of a substance. q = (specific heat)  (grams of substance) T. Be careful of the sign of q. Chapter 5

  16. Calorimetry Heat Capacity and Specific Heat Chapter 5

  17. Calorimetry Constant-Pressure Calorimetry Atmospheric pressure is constant! DH = qP qrxn = -qsoln = -(specific heat of solution)  (grams of solution) DT. Chapter 5

  18. Calorimetry Bomb Calorimetry (Constant-Volume Calorimetry) Reaction carried out under constant volume. Use a bomb calorimeter. Usually study combustion. Chapter 5

  19. Calorimetry Bomb Calorimetry (Constant-Volume Calorimetry) qrxn = -CcalorimeterT. Chapter 5

  20. Hess’s Law • Hess’s law: if a reaction is carried out in a number of steps, H for the overall reaction is the sum of H for each individual step. • For example: CH4(g) + 2O2(g)  CO2(g) + 2H2O(g) H = -802 kJ 2H2O(g)  2H2O(l) H = -88 kJ CH4(g) + 2O2(g)  CO2(g) + 2H2O(l) H = -890 kJ Chapter 5

  21. Hess’s Law In the above enthalpy diagram note that H1 = H2 + H3 Chapter 5

  22. Enthalpies of Formation • If 1 mol of compound is formed from its constituent elements, then the enthalpy change for the reaction is called the enthalpy of formation, Hof . • Standard conditions (standard state): 1 atm and 25 oC (298 K). • Standard enthalpy, Ho, is the enthalpy measured when everything is in its standard state. • Standard enthalpy of formation: 1 mol of compound is formed from substances in their standard states. • If there is more than one state for a substance under standard conditions, the more stable one is used. Chapter 5

  23. Enthalpies of Formation • Standard enthalpy of formation of the most stable form of an element is zero. Chapter 5

  24. Enthalpies of Formation Using Enthalpies of Formation to Calculate Enthalpies of Reaction We use Hess’ Law to calculate enthalpies of a reaction from enthalpies of formation. Hrxn = H1 + H2 + H3 Chapter 5

  25. Enthalpies of Formation Using Enthalpies of Formation to Calculate Enthalpies of Reaction For a reaction: Chapter 5

  26. Foods and Fuels • Foods • Fuel value = energy released when 1 g of substance is burned. • 1 nutritional Calorie, 1 Cal = 1000 cal = 1 kcal. • Energy in our bodies comes from carbohydrates and fats (mostly). • Intestines: carbohydrates converted into glucose: • C6H12O6 + 6O2 6CO2 + 6H2O, DH = -2816 kJ • Fats break down as follows: • 2C57H110O6 + 163O2 114CO2 + 110H2O, DH = -75,520 kJ • Fats: contain more energy; are not water soluble, so are good for energy storage. Chapter 5

  27. Foods and Fuels Foods Chapter 5

  28. Foods and Fuels Fuels U.S.: 1.0 x 106 kJ of fuel per day. Most from petroleum and natural gas. Remainder from coal, nuclear, and hydroelectric. Fossil fuels are not renewable. Chapter 5

  29. Foods and Fuels Fuels Fuel value = energy released when 1 g of substance is burned. Hydrogen has great potential as a fuel with a fuel value of 142 kJ/g. Chapter 5

  30. Thermochemistry End of Chapter 5 Chapter 5