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Chapter 6Thermochemistry
Chapter 6      Table of Contents      6.1             The Nature of Energy      6.2             Enthalpy and Calorimetry  ...
Section 6.1      The Nature of Energy      Energy       • Capacity to do work or to produce heat.       • That which is ne...
Section 6.1      The Nature of Energy      Energy       •         Potential energy – energy due to position               ...
Section 6.1      The Nature of Energy      Initial Position       •         In the initial position, ball A has a higher  ...
Section 6.1      The Nature of Energy      Final Position       •         After A has rolled down the hill, the potential ...
Section 6.1      The Nature of Energy      Energy       • Heat involves the transfer of energy between         two objects...
Section 6.1      The Nature of Energy      Chemical Energy       •         System – part of the universe on which we      ...
Section 6.1      The Nature of Energy      Chemical Energy       • Endothermic Reaction:          Heat flow is into a sys...
Section 6.1      The Nature of Energy                         Concept Check                  Is the freezing of water an e...
Section 6.1      The Nature of Energy                         Concept Check                     Classify each process as e...
Section 6.1      The Nature of Energy                         Concept Check                  For each of the following, de...
Section 6.1      The Nature of Energy                         Concept Check                  Hydrogen gas and oxygen gas r...
Section 6.1      The Nature of Energy      Internal Energy       •         Law of conservation of energy is often         ...
Section 6.1      The Nature of Energy      Work vs. Energy Flow                                                    Return ...
Section 6.1      The Nature of Energy      Internal Energy       • Sign reflects the system’s point of view.       • Endot...
Section 6.1      The Nature of Energy      Internal Energy       • Sign reflects the system’s point of view.       • Syste...
Section 6.1      The Nature of Energy      Work       •         Work = P × A × Δh = PΔV                      P is pressur...
Section 6.1      The Nature of Energy      Work       •         For an expanding gas, ΔV is a positive                 qua...
Section 6.1      The Nature of Energy                         Exercise                   Which of the following performs m...
Section 6.1      The Nature of Energy                       Concept Check                      Determine the sign of ∆E fo...
Section 6.2      Enthalpy and Calorimetry      Change in Enthalpy       •         State function       •         ΔH = q at...
Section 6.2      Enthalpy and Calorimetry                         Exercise                   Consider the combustion of pr...
Section 6.2      Enthalpy and Calorimetry      Calorimetry       • Science of measuring heat       • Specific heat capacit...
Section 6.2      Enthalpy and Calorimetry      Calorimetry       •         If two reactants at the same temperature       ...
Section 6.2      Enthalpy and Calorimetry      A Coffee–Cup      Calorimeter Made of      Two Styrofoam      Cups         ...
Section 6.2      Enthalpy and Calorimetry      Calorimetry              •        Energy released (heat) = s × m × ΔT      ...
Section 6.2      Enthalpy and Calorimetry                         Concept Check                  A 100.0 g sample of water...
Section 6.2      Enthalpy and Calorimetry                         Concept Check                  A 100.0 g sample of water...
Section 6.2      Enthalpy and Calorimetry         Concept Check                  You have a Styrofoam cup with 50.0 g of w...
Section 6.3      Hess’s Law       •         In going from a particular set of reactants                 to a particular se...
Section 6.3      Hess’s Law      N2(g) + 2O2(g) → 2NO2(g)                              ΔH1 = 68 kJ       •         This re...
Section 6.3      Hess’s Law      The Principle of Hess’s Law                                                    Return to ...
Section 6.3      Hess’s Law                                                    Return to TOCCopyright © Cengage Learning. ...
Section 6.3      Hess’s Law      Characteristics of Enthalpy Changes       •         If a reaction is reversed, the sign o...
Section 6.3      Hess’s Law      Example       •         Consider the following data:                                     ...
Section 6.3      Hess’s Law      Problem-Solving Strategy       •         Work backward from the required                 ...
Section 6.3      Hess’s Law      Example       •         Reverse the two reactions:                  1                    ...
Section 6.3      Hess’s Law      Example       •         Multiply reactions to give the correct numbers                 of...
Section 6.3      Hess’s Law      Example       •         Final reactions:                 2 N2 ( g ) + 6 H2 ( g )  4 NH3...
Section 6.4      Standard Enthalpies of Formation      Standard Enthalpy of Formation (ΔHf°)       •         Change in ent...
Section 6.4      Standard Enthalpies of Formation      Conventional Definitions of Standard States       • For a Compound ...
Section 6.4      Standard Enthalpies of Formation           A Schematic Diagram of the Energy Changes for the Reaction    ...
Section 6.4      Standard Enthalpies of Formation      Problem-Solving Strategy: Enthalpy Calculations       1. When a rea...
Section 6.4      Standard Enthalpies of Formation      Problem-Solving Strategy: Enthalpy Calculations       3. The change...
Section 6.4      Standard Enthalpies of Formation                         Exercise                         Calculate ∆H° f...
Section 6.5      Present Sources of Energy       • Fossil Fuels          Petroleum, Natural Gas, and Coal       • Wood   ...
Section 6.5      Present Sources of Energy      Energy Sources Used in the United States                                  ...
Section 6.5      Present Sources of Energy      The Earth’s Atmosphere       •         Transparent to visible light from t...
Section 6.5      Present Sources of Energy      The Earth’s Atmosphere                                                    ...
Section 6.6      New Energy Sources       •       Coal Conversion       •       Hydrogen as a Fuel       •       Other Ene...
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thermodynamics

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  • Exothermic process because you must remove energy in order to slow the molecules down to form a solid.
  • Exothermic (heat energy leaves your hand and moves to the ice) Endothermic (heat energy flows into the ice) Endothermic (heat energy flows into the water to boil it) Exothermic (heat energy leaves to condense the water from a gas to a liquid) Endothermic (heat energy flows into the ice cream to melt it)
  • System – methane and oxygen to produce carbon dioxide and water; Surroundings – everything else around it; Direction of energy transfer – energy transfers from the system to the surroundings (exothermic) System – water drops; Surroundings – everything else around it; Direction of energy transfer – energy transfers from the surroundings (your body) to the system (water drops) (endothermic)
  • Water is lower in energy because a lot of energy was released in the process when hydrogen and oxygen gases reacted.
  • Answer to questions in animation: Work ( w ) is negative. Heat ( q ) is negative. Internal energy ( Δ E ) changes.
  • They both perform the same amount of work. w = - P Δ V a) w = -(2 atm)(4.0-1.0) = -6 L·atm b) w = -(3 atm)(3.0-1.0) = -6 L·atm
  • a) q is positive for endothermic processes and w is negative when system does work on surroundings; first condition – Δ E is negative; second condition – Δ E is positive b) q is negative for exothermic processes and w is positive when surroundings does work on system; first condition – Δ E is positive; second condition – Δ E is negative
  • (5.00 g C 3 H 8 )(1 mol / 44.094 g C 3 H 8 )(-2221 kJ / mol C 3 H 8 ) Δ H = -252 kJ
  • The correct answer is b).
  • The correct answer is c). The final temperature of the water is 23 °C. - (100.0 g)(4.18 J/°C·g)(T f – 90.) = (500.0 g)(4.18 J/°C·g)(T f – 10.) T f = 23°C
  • The correct answer is c). The final temperature of the water is 18 °C. - (50.0 g)(0.45 J/°C·g)(T f – 90.) = (50.0 g)(4.18 J/°C·g)(T f – 10.) T f = 18°C
  • [2(–470) + 0] – [0 + 2(–286)] = –368 kJ Δ H = –368 kJ
  • Transcript of "thermodynamics"

    1. 1. Chapter 6Thermochemistry
    2. 2. Chapter 6 Table of Contents 6.1 The Nature of Energy 6.2 Enthalpy and Calorimetry 6.3 Hess’s Law 6.4 Standard Enthalpies of Formation 6.5 Present Sources of Energy 6.6 New Energy SourcesCopyright © Cengage Learning. All rights reserved 2
    3. 3. Section 6.1 The Nature of Energy Energy • Capacity to do work or to produce heat. • That which is needed to oppose natural attractions. • Law of conservation of energy – energy can be converted from one form to another but can be neither created nor destroyed.  The total energy content of the universe is constant. Return to TOCCopyright © Cengage Learning. All rights reserved 3
    4. 4. Section 6.1 The Nature of Energy Energy • Potential energy – energy due to position or composition. • Kinetic energy – energy due to motion of the object and depends on the mass of the object and its velocity. Return to TOCCopyright © Cengage Learning. All rights reserved 4
    5. 5. Section 6.1 The Nature of Energy Initial Position • In the initial position, ball A has a higher potential energy than ball B. Return to TOCCopyright © Cengage Learning. All rights reserved 5
    6. 6. Section 6.1 The Nature of Energy Final Position • After A has rolled down the hill, the potential energy lost by A has been converted to random motions of the components of the hill (frictional heating) and to the increase in the potential energy of B. Return to TOCCopyright © Cengage Learning. All rights reserved 6
    7. 7. Section 6.1 The Nature of Energy Energy • Heat involves the transfer of energy between two objects due to a temperature difference. • Work – force acting over a distance. • Energy is a state function; work and heat are not:  State Function – property that does not depend in any way on the system’s past or future (only depends on present state). Return to TOCCopyright © Cengage Learning. All rights reserved 7
    8. 8. Section 6.1 The Nature of Energy Chemical Energy • System – part of the universe on which we wish to focus attention. • Surroundings – include everything else in the universe. Return to TOCCopyright © Cengage Learning. All rights reserved 8
    9. 9. Section 6.1 The Nature of Energy Chemical Energy • Endothermic Reaction:  Heat flow is into a system.  Absorb energy from the surroundings. • Exothermic Reaction:  Energy flows out of the system. • Energy gained by the surroundings must be equal to the energy lost by the system. Return to TOCCopyright © Cengage Learning. All rights reserved 9
    10. 10. Section 6.1 The Nature of Energy Concept Check Is the freezing of water an endothermic or exothermic process? Explain. Return to TOCCopyright © Cengage Learning. All rights reserved 10
    11. 11. Section 6.1 The Nature of Energy Concept Check Classify each process as exothermic or endothermic. Explain. The system is underlined in each example. Exo a) Your hand gets cold when you touch ice. Endo b) The ice gets warmer when you touch it. Endo c) Water boils in a kettle being heated on a stove. Exo d) Water vapor condenses on a cold pipe. Endo e) Ice cream melts. Return to TOCCopyright © Cengage Learning. All rights reserved 11
    12. 12. Section 6.1 The Nature of Energy Concept Check For each of the following, define a system and its surroundings and give the direction of energy transfer. a) Methane is burning in a Bunsen burner in a laboratory. b) Water drops, sitting on your skin after swimming, evaporate. Return to TOCCopyright © Cengage Learning. All rights reserved 12
    13. 13. Section 6.1 The Nature of Energy Concept Check Hydrogen gas and oxygen gas react violently to form water. Explain.  Which is lower in energy: a mixture of hydrogen and oxygen gases, or water? Return to TOCCopyright © Cengage Learning. All rights reserved 13
    14. 14. Section 6.1 The Nature of Energy Internal Energy • Law of conservation of energy is often called the first law of thermodynamics. • Internal energy E of a system is the sum of the kinetic and potential energies of all the “particles” in the system. • To change the internal energy of a system: ΔE = q + w q represents heat w represents work Return to TOCCopyright © Cengage Learning. All rights reserved 14
    15. 15. Section 6.1 The Nature of Energy Work vs. Energy Flow Return to TOCCopyright © Cengage Learning. All rights reserved 15
    16. 16. Section 6.1 The Nature of Energy Internal Energy • Sign reflects the system’s point of view. • Endothermic Process:  q is positive • Exothermic Process:  q is negative Return to TOCCopyright © Cengage Learning. All rights reserved 16
    17. 17. Section 6.1 The Nature of Energy Internal Energy • Sign reflects the system’s point of view. • System does work on surroundings:  w is negative • Surroundings do work on the system:  w is positive Return to TOCCopyright © Cengage Learning. All rights reserved 17
    18. 18. Section 6.1 The Nature of Energy Work • Work = P × A × Δh = PΔV  P is pressure.  A is area.  Δh is the piston moving a distance.  ΔV is the change in volume. Return to TOCCopyright © Cengage Learning. All rights reserved 18
    19. 19. Section 6.1 The Nature of Energy Work • For an expanding gas, ΔV is a positive quantity because the volume is increasing. Thus ΔV and w must have opposite signs: w = –PΔV • To convert between L·atm and Joules, use 1 L·atm = 101.3 J. Return to TOCCopyright © Cengage Learning. All rights reserved 19
    20. 20. Section 6.1 The Nature of Energy Exercise Which of the following performs more work? a) A gas expanding against a pressure of 2 atm from 1.0 L to 4.0 L. b) A gas expanding against a pressure of 3 atm from 1.0 L to 3.0 L. They perform the same amount of work. Return to TOCCopyright © Cengage Learning. All rights reserved 20
    21. 21. Section 6.1 The Nature of Energy Concept Check Determine the sign of ∆E for each of the following with the listed conditions: a) An endothermic process that performs work.  |work| > |heat| Δ E = negative  |work| < |heat| Δ E = positive b) Work is done on a gas and the process is exothermic.  |work| > |heat| Δ E = positive  |work| < |heat| Δ E = negative Return to TOCCopyright © Cengage Learning. All rights reserved 21
    22. 22. Section 6.2 Enthalpy and Calorimetry Change in Enthalpy • State function • ΔH = q at constant pressure • ΔH = Hproducts – Hreactants Return to TOCCopyright © Cengage Learning. All rights reserved 22
    23. 23. Section 6.2 Enthalpy and Calorimetry Exercise Consider the combustion of propane: C3H8(g) + 5O2(g) → 3CO2(g) + 4H2O(l) ΔH = –2221 kJ Assume that all of the heat comes from the combustion of propane. Calculate ΔH in which 5.00 g of propane is burned in excess oxygen at constant pressure. –252 kJ Return to TOCCopyright © Cengage Learning. All rights reserved 23
    24. 24. Section 6.2 Enthalpy and Calorimetry Calorimetry • Science of measuring heat • Specific heat capacity:  The energy required to raise the temperature of one gram of a substance by one degree Celsius. • Molar heat capacity:  The energy required to raise the temperature of one mole of substance by one degree Celsius. Return to TOCCopyright © Cengage Learning. All rights reserved 24
    25. 25. Section 6.2 Enthalpy and Calorimetry Calorimetry • If two reactants at the same temperature are mixed and the resulting solution gets warmer, this means the reaction taking place is exothermic. • An endothermic reaction cools the solution. Return to TOCCopyright © Cengage Learning. All rights reserved 25
    26. 26. Section 6.2 Enthalpy and Calorimetry A Coffee–Cup Calorimeter Made of Two Styrofoam Cups Return to TOCCopyright © Cengage Learning. All rights reserved 26
    27. 27. Section 6.2 Enthalpy and Calorimetry Calorimetry • Energy released (heat) = s × m × ΔT s = specific heat capacity (J/°C·g) m = mass (g) ΔT = change in temperature (°C) Return to TOCCopyright © Cengage Learning. All rights reserved 27
    28. 28. Section 6.2 Enthalpy and Calorimetry Concept Check A 100.0 g sample of water at 90°C is added to a 100.0 g sample of water at 10°C. The final temperature of the water is: a) Between 50°C and 90°C b) 50°C c) Between 10°C and 50°C Return to TOCCopyright © Cengage Learning. All rights reserved 28
    29. 29. Section 6.2 Enthalpy and Calorimetry Concept Check A 100.0 g sample of water at 90.°C is added to a 500.0 g sample of water at 10.°C. The final temperature of the water is: a) Between 50°C and 90°C b) 50°C c) Between 10°C and 50°C Calculate the final temperature of the water. 23°C Return to TOCCopyright © Cengage Learning. All rights reserved 29
    30. 30. Section 6.2 Enthalpy and Calorimetry Concept Check You have a Styrofoam cup with 50.0 g of water at 10.°C. You add a 50.0 g iron ball at 90.°C to the water. (sH2O = 4.18 J/°C·g and sFe = 0.45 J/°C·g) The final temperature of the water is: a) Between 50°C and 90°C b) 50°C c) Between 10°C and 50°C Calculate the final temperature of the water. 18°C Return to TOCCopyright © Cengage Learning. All rights reserved 30
    31. 31. Section 6.3 Hess’s Law • In going from a particular set of reactants to a particular set of products, the change in enthalpy is the same whether the reaction takes place in one step or in a series of steps. Return to TOCCopyright © Cengage Learning. All rights reserved 31
    32. 32. Section 6.3 Hess’s Law N2(g) + 2O2(g) → 2NO2(g) ΔH1 = 68 kJ • This reaction also can be carried out in two distinct steps, with enthalpy changes designated by ΔH2 and ΔH3. N2(g) + O2(g) → 2NO(g) ΔH2 = 180 kJ 2NO(g) + O2(g) → 2NO2(g) ΔH3 = – 112 kJ N2(g) + 2O2(g) → 2NO2(g) ΔH2 + ΔH3 = 68 kJ ΔH1 = ΔH2 + ΔH3 = 68 kJ Return to TOCCopyright © Cengage Learning. All rights reserved 32
    33. 33. Section 6.3 Hess’s Law The Principle of Hess’s Law Return to TOCCopyright © Cengage Learning. All rights reserved 33
    34. 34. Section 6.3 Hess’s Law Return to TOCCopyright © Cengage Learning. All rights reserved 34
    35. 35. Section 6.3 Hess’s Law Characteristics of Enthalpy Changes • If a reaction is reversed, the sign of ΔH is also reversed. • The magnitude of ΔH is directly proportional to the quantities of reactants and products in a reaction. If the coefficients in a balanced reaction are multiplied by an integer, the value of ΔH is multiplied by the same integer. Return to TOCCopyright © Cengage Learning. All rights reserved 35
    36. 36. Section 6.3 Hess’s Law Example • Consider the following data: 1 3 NH3 (g )  N2 ( g ) + H2 (g ) → ∆H = 46 kJ 2 2 2 H2 ( g ) + O2 ( g )  2 H2O( g ) → ∆H = − 484 kJ • Calculate ΔH for the reaction 2 N2 ( g ) + 6 H2O( g )  3 O 2 ( g ) + 4 NH3 ( g ) → Return to TOCCopyright © Cengage Learning. All rights reserved 36
    37. 37. Section 6.3 Hess’s Law Problem-Solving Strategy • Work backward from the required reaction, using the reactants and products to decide how to manipulate the other given reactions at your disposal. • Reverse any reactions as needed to give the required reactants and products. • Multiply reactions to give the correct numbers of reactants and products. Return to TOCCopyright © Cengage Learning. All rights reserved 37
    38. 38. Section 6.3 Hess’s Law Example • Reverse the two reactions: 1 3 N2 (g ) + H2 ( g )  NH3 (g ) → ∆H = − 46 kJ 2 2 2 H2O( g )  2 H2 ( g ) + O2 ( g ) → ∆H = +484 kJ • Desired reaction: 2 N2 ( g ) + 6 H2O( g )  3 O 2 ( g ) + 4 NH3 ( g ) → Return to TOCCopyright © Cengage Learning. All rights reserved 38
    39. 39. Section 6.3 Hess’s Law Example • Multiply reactions to give the correct numbers of reactants and products: 1 3 4( 2 N2 (g ) + H2 ( g )  NH3 (g ) ) → 4( ∆H = − 46 kJ ) 2 3( 2 H2O(g )  2 H2 (g ) + O2 (g )) 3(∆H = +484 kJ ) → • Desired reaction: 2 N2 ( g ) + 6 H2O( g )  3 O2 ( g ) + 4 NH3 ( g ) → Return to TOCCopyright © Cengage Learning. All rights reserved 39
    40. 40. Section 6.3 Hess’s Law Example • Final reactions: 2 N2 ( g ) + 6 H2 ( g )  4 NH3 ( g ) → ∆H = − 184 kJ 6 H2O( g )  6 H2 ( g ) + 3 O 2 ( g ) → ∆H = +1452 kJ • Desired reaction: 2 N2 ( g ) + 6 H2O( g )  3 O 2 ( g ) + 4 NH3 ( g ) → ΔH = +1268 kJ Return to TOCCopyright © Cengage Learning. All rights reserved 40
    41. 41. Section 6.4 Standard Enthalpies of Formation Standard Enthalpy of Formation (ΔHf°) • Change in enthalpy that accompanies the formation of one mole of a compound from its elements with all substances in their standard states. Return to TOCCopyright © Cengage Learning. All rights reserved 41
    42. 42. Section 6.4 Standard Enthalpies of Formation Conventional Definitions of Standard States • For a Compound  For a gas, pressure is exactly 1 atm.  For a solution, concentration is exactly 1 M.  Pure substance (liquid or solid) • For an Element  The form [N2(g), K(s)] in which it exists at 1 atm and 25°C.  Heat of formation is zero. Return to TOCCopyright © Cengage Learning. All rights reserved 42
    43. 43. Section 6.4 Standard Enthalpies of Formation A Schematic Diagram of the Energy Changes for the Reaction CH4(g) + 2O2(g) → CO2(g) + 2H2O(l) ΔH°reaction = –(–75 kJ) + 0 + (–394 kJ) + (–572 kJ) = –891 kJ Return to TOCCopyright © Cengage Learning. All rights reserved 43
    44. 44. Section 6.4 Standard Enthalpies of Formation Problem-Solving Strategy: Enthalpy Calculations 1. When a reaction is reversed, the magnitude of ΔH remains the same, but its sign changes. 2. When the balanced equation for a reaction is multiplied by an integer, the value of ΔH for that reaction must be multiplied by the same integer. Return to TOCCopyright © Cengage Learning. All rights reserved 44
    45. 45. Section 6.4 Standard Enthalpies of Formation Problem-Solving Strategy: Enthalpy Calculations 3. The change in enthalpy for a given reaction can be calculated from the enthalpies of formation of the reactants and products: ∆H°rxn = Σnp∆Hf°(products) - Σnr∆Hf°(reactants) 4. Elements in their standard states are not included in the ΔHreaction calculations because ΔHf° for an element in its standard state is zero. Return to TOCCopyright © Cengage Learning. All rights reserved 45
    46. 46. Section 6.4 Standard Enthalpies of Formation Exercise Calculate ∆H° for the following reaction: 2Na(s) + 2H2O(l) → 2NaOH(aq) + H2(g) Given the following information: ∆Hf° (kJ/mol) Na(s) 0 H2O(l) –286 NaOH(aq) –470 H2(g) 0 ∆H° = –368 kJ Return to TOCCopyright © Cengage Learning. All rights reserved 46
    47. 47. Section 6.5 Present Sources of Energy • Fossil Fuels  Petroleum, Natural Gas, and Coal • Wood • Hydro • Nuclear Return to TOCCopyright © Cengage Learning. All rights reserved 47
    48. 48. Section 6.5 Present Sources of Energy Energy Sources Used in the United States Return to TOCCopyright © Cengage Learning. All rights reserved 48
    49. 49. Section 6.5 Present Sources of Energy The Earth’s Atmosphere • Transparent to visible light from the sun. • Visible light strikes the Earth, and part of it is changed to infrared radiation. • Infrared radiation from Earth’s surface is strongly absorbed by CO2, H2O, and other molecules present in smaller amounts in atmosphere. • Atmosphere traps some of the energy and keeps the Earth warmer than it would otherwise be. Return to TOCCopyright © Cengage Learning. All rights reserved 49
    50. 50. Section 6.5 Present Sources of Energy The Earth’s Atmosphere Return to TOCCopyright © Cengage Learning. All rights reserved 50
    51. 51. Section 6.6 New Energy Sources • Coal Conversion • Hydrogen as a Fuel • Other Energy Alternatives  Oil shale  Ethanol  Methanol  Seed oil Return to TOCCopyright © Cengage Learning. All rights reserved 51
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