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Chapter 6

Chapter 6. Principles of Reactivity: Energy and Chemical Reactions. Thermite Reaction. Terminology. Energy capacity to do work Kinetic Energy energy that something has because it is moving Potential Energy energy that something has because of its position or its chemical bonding.

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Chapter 6

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  1. Chapter 6 Principles of Reactivity: Energy and Chemical Reactions Dr. S. M. Condren

  2. Thermite Reaction Dr. S. M. Condren

  3. Terminology Energy • capacity to do work Kinetic Energy • energy that something has because it is moving Potential Energy • energy that something has because of its position or its chemical bonding Dr. S. M. Condren

  4. Kinetic Energy Dr. S. M. Condren

  5. Chemical Potential Energy Dr. S. M. Condren

  6. Chemical Potential Energy Dr. S. M. Condren

  7. Internal Energy • The sum of the individual energies of all nanoscale particles (atoms, ions, or molecules) in that sample. • E = 1/2mc2 • The total internal energy of a sample of matter depends on temperature, the type of particles, and how many of them there are in the sample. Dr. S. M. Condren

  8. Energy Units • calorie - energy required to heat 1-g of water 1oC • Calorie - unit of food energy; • 1 Cal = 1-kcal = 1000-cal • Joule - 1-cal = 4.184 J = 1-kg*m2/sec2 Dr. S. M. Condren

  9. Law of Conservation of Energy • energy can neither be created nor destroyed • the total amount of energy in the universe is a constant • energy can be transformed from one form to another Dr. S. M. Condren

  10. First Law of Thermodynamics • the amount of heat transferred into a system plus the amount of work done on the system must result in a corresponding increase of internal energy in the system Dr. S. M. Condren

  11. Thermochemistry Terminology system => that part of the universe under investigation surroundings => the rest of the universe universe = system + surroundings Dr. S. M. Condren

  12. System and Surroundings • SYSTEM • The object under study • SURROUNDINGS • Everything outside the system Dr. S. M. Condren

  13. Announcement Learn@UW will be unavailable on Thursday July 19 from 5:00am until 12:00 noon. Dr. S. M. Condren

  14. Energy & Chemistry 2 H2(g) + O2(g) --> 2 H2O(g) + heat and light This can be set up to provide ELECTRIC ENERGY in a fuel cell. Oxidation: 2 H2 ---> 4 H+ + 4 e- Reduction: 4 e- + O2 + 2 H2O ---> 4 OH- H2/O2 Fuel Cell Energy, page 288 Dr. S. M. Condren

  15. Energy & Chemistry ENERGY is the capacity to do work or transfer heat. HEAT is the form of energy that flows between 2 objects because of their difference in temperature. Other forms of energy — • light • electrical • kinetic and potential Dr. S. M. Condren

  16. Potential Energy in the Atomic Scale • Positive and negative particles (ions) attract one another. • Two atoms can bond • As the particles attract they have a lower potential energy NaCl — composed of Na+ and Cl- ions. http://mrsec.wisc.edu/Edetc/pmk/NaCl_alt.html Dr. S. M. Condren

  17. Internal Energy (E) PE + KE = Internal energy (E or U) Int. E of a chemical system depends on • number of particles • type of particles • temperature Dr. S. M. Condren

  18. Energy Transfer Energy is always transferred from the hotter to the cooler sample Heat – the energy that flows into or out of a system because of a difference in temperature between the thermodynamic system and its surroundings Dr. S. M. Condren

  19. Thermochemistry Terminology state properties => properties which depend only on the initial and final states => properties which are path independent non-state properties => properties which are path dependent state properties => E non-state properties => q & w Dr. S. M. Condren

  20. Thermochemistry Terminology exothermic - reaction that gives off energy endothermic - reaction that absorbs energy chemical energy - energy associated with a chemical reaction thermochemistry - the quantitative study of the heat changes accompanying chemical reactions thermodynamics - the study of energy and its transformations Dr. S. M. Condren

  21. Exothermic Reaction First-Aid Hotpacks, containing either calcium chloride or magnesium sulfate, plus water Dr. S. M. Condren

  22. Endothermic Reaction First-aid cold packs, containing ammonium nitrate and water in separate inner pouches Dr. S. M. Condren

  23. Enthalpy • heat at constant pressure qp = DH = Hproducts - Hreactants Exothermic Reaction DH = (Hproducts - Hreactants) < 0 H2O(l) -----> H2O(s)DH < 0 Endothermic Reaction DH = (Hproducts - Hreactants) > 0 H2O(l) -----> H2O(g)DH > 0 Dr. S. M. Condren

  24. Enthalpy H = E + PV DH = DE + PDV DE = DH – PDV Dr. S. M. Condren

  25. First Law of Thermodynamics heat => q internal energy => E internal energy change =>DE work => w = - P*DV DE = q + w Dr. S. M. Condren

  26. Specific Heat-Specific Heat Capacity • the amount of heat necessary to raise the temperature of 1 gram of the substance 1oC • independent of mass • substance dependent • s.h. • Specific Heat of Water = 4.184 J/goC Dr. S. M. Condren

  27. Heat q = m * s.h. * Dt where q => heat, J m => mass, g s.h. => specific heat, J/g*oC Dt = change in temperature, oC, (always tf – ti) Dr. S. M. Condren

  28. Molar Heat Capacity • the heat necessary to raise the temperature of one mole of substance by 1oC • substance dependent • C Dr. S. M. Condren

  29. Heat Capacity • the heat necessary to raise the temperature 1oC • mass dependent • substance dependent • C Dr. S. M. Condren

  30. Heat Capacity C = m X s.h. where C => heat capacity, J/oC m => mass, g s.h. => specific heat, J/goC Dr. S. M. Condren

  31. Plotted are graphs of heat absorbed versus temperature for two systems. Which system has the larger heat capacity? A, B Dr. S. M. Condren

  32. Heat Transfer qlost = - qgained (m X s.h. X Dt)lost = - (m X s.h. X Dt)gained Dr. S. M. Condren

  33. Heat Transfer Dr. S. M. Condren

  34. EXAMPLEIf 100. g of iron at 100.0oC is placed in 200. g of water at 20.0oC in an insulated container, what will the temperature, oC, of the iron and water when both are at the same temperature? The specific heat of iron is 0.106 cal/goC. (100.g*0.106cal/goC*(Tf - 100.)oC) = qlost - qgained = (200.g*1.00cal/goC*(Tf - 20.0)oC) 10.6(Tf - 100.oC) = - 200.(Tf - 20.0oC) 10.6Tf - 1060oC = - 200.Tf + 4000oC (10.6 + 200.)Tf = (1060 + 4000)oC Tf = (5060/211.)oC = 24.0oC Dr. S. M. Condren

  35. Melting of Ice http://mrsec.wisc.edu/Edetc/pmk/ice.html Dr. S. M. Condren

  36. EXAMPLE: How much heat is required to heat 10.0 g of ice at -15.0oC to steam at 127.0oC? q = DHice + DHfusion + DHwater + DHboil. + DHsteam q = DHice + DHfusion + DHwater + DHboil. + DHsteam Dr. S. M. Condren

  37. Heat Transfer Dr. S. M. Condren

  38. EXAMPLE: How much heat is required to heat 10.0 g of ice at -15.0oC to steam at 127.0oC?q = DHice + DHfusion + DHwater + DHboil. + DHsteam q = (10.0g*2.09J/goC*((0.0 – (-15.0))oC)) { specific heat of ice Mass of the ice Temperature change Dr. S. M. Condren

  39. EXAMPLE: How much heat is required to heat 10.0 g of ice at -15.0oC to steam at 127.0oC?q = DHice + DHfusion + DHwater + DHboil. + DHsteam q = (10.0g*2.09J/goC*15.0oC) + (10.0g*333J/g) Melting of ice occurs at a constant temperature Mass of ice Heat of fusion Dr. S. M. Condren

  40. EXAMPLE: How much heat is required to heat 10.0 g of ice at -15.0oC to steam at 127.0oC?q = DHice + DHfusion + DHwater + DHboil. + DHsteam q = (10.0g*2.09J/goC*15.0oC) + (10.0g*333J/g) + (10.0g*4.18J/goC*((100.0-0.00)oC)) Mass of water Specific heat of liquid water Temperature change of the liquid water Dr. S. M. Condren

  41. EXAMPLE: How much heat is required to heat 10.0 g of ice at -15.0oC to steam at 127.0oC?q = DHice + DHfusion + DHwater + DHboil. + DHsteam q = (10.0g*2.09J/goC*15.0oC) + (10.0g*333J/g) + (10.0g*4.18J/goC*100.0oC) + (10.0g*2260J/g) Boiling of water occurs at a constant temperature Mass of water Heat of vaporization Dr. S. M. Condren

  42. EXAMPLE: How much heat is required to heat 10.0 g of ice at -15.0oC to steam at 127.0oC?q = DHice + DHfusion + DHwater + DHboil. + DHsteam q = (10.0g*2.09J/goC*15.0oC) + (10.0g*333J/g) + (10.0g*4.18J/goC*100.0oC) + (10.0g*2260J/g) + (10.0g*2.03J/goC*((127.0-100.0)oC)) Specific heatof steam Temperature change for the steam Mass of steam Dr. S. M. Condren

  43. EXAMPLE: How much heat is required to heat 10.0 g of ice at -15.0oC to steam at 127.0oC?q = DHice + DHfusion + DHwater + DHboil. + DHsteam q = (10.0g*2.09J/goC*15.0oC) + (10.0g*333J/g) + (10.0g*4.18J/goC*100.0oC) + (10.0g*2260J/g) + (10.0g*2.03J/goC*27.0oC) q = (314 )J + 3.33X103 + 4.18X103 + 2.26X104 + 548 = 30.96 kJ Dr. S. M. Condren

  44. Spreadsheet of Previous Problem Dr. S. M. Condren

  45. Coffee Cup Calorimeter Dr. S. M. Condren

  46. Bomb Calorimeter Dr. S. M. Condren

  47. EXAMPLE A 1.000g sample of a particular compound produced 11.0 kJ of heat. The temperature of the calorimeter and 3000 g of water was raised 0.629oC. How much heat is gained by the calorimeter? heat gained = - heat lost heatcalorimeter + heatwater = heatreaction heatcalorimeter = heatreaction - heatwater Dr. S. M. Condren

  48. EXAMPLE A 1.000g sample of a particular compound produced 11.0 kJ of heat. The temperature of the calorimeter and 3000. g of water was raised 0.629oC. How much heat is gained by the calorimeter? heatcalorimeter = heatreaction - heatwater heat = 11.0 kJ - ((3.000kg)(0.629oC)(4.184kJ/kgoC)) = 3.10 kJ Dr. S. M. Condren

  49. Example What is the mass of water equivalent of the heat absorbed by the calorimeter? #g = (3.10 kJ/0.629oC)(1.00kg*oC/4.184kJ) = 6.47 x 102 g Dr. S. M. Condren

  50. Example A 1.000 g sample of ethanol was burned in the sealed bomb calorimeter described above. The temperature of the water rose from 24.284oC to 26.225oC. Determine the heat for the reaction. m = (3000. + 647)g H2O q = m X s.h. X Dt = (3647g)(4.184J/goC)(1.941oC) = 29.61 kJ Dr. S. M. Condren

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