New chm 151_unit_9_power_points

download New chm 151_unit_9_power_points

of 83

  • date post

    28-Aug-2014
  • Category

    Documents

  • view

    339
  • download

    0

Embed Size (px)

description

 

Transcript of New chm 151_unit_9_power_points

  • THERMOCHEMISTRY 6-1 TERMINOLOGY ENTHALPY CHANGES BOND ENERGIES Chapter 6.1-6.6 & 9.4 Silberberg
  • Chapter 6 Thermochemistry: Energy Flow and Chemical Change 6-2
  • Thermochemistry: Energy Flow and Chemical Change 6.1 Forms of Energy and Their Interconversion 6.2 Enthalpy: Chemical Change at Constant Pressure 6.3 Calorimetry: Measuring the Heat of a Chemical or Physical Change 6.4 Stoichiometry of Thermochemical Equations 6.5 Hesss Law: Finding DH of Any Reaction 6.6 Standard Enthalpies of Reaction (DH 6-3 rxn)
  • Models of Chemical Bonding 9.4 Bond Energy and Chemical Change 6-4
  • Goals & Objectives See the following Learning Objectives on pages 253 and 356. Understand these Concepts: 6.1-4, 6-12; 9.11-12 Master these Skills: 6.1, 3-6; 9.5 6-5
  • Thermochemistry Thermochemistry is a branch of thermodynamics that deals with the heat involved with chemical and physical changes. When energy is transferred from one object to another it appears as work and/or heat. 6-6
  • Transfer and Interconversion of Energy Thermodynamics is the study of energy and its transformations. Thermochemistry is a branch of thermodynamics that deals with the heat involved in chemical and physical changes. When energy is transferred from one object to another, it appears as work and heat. 6-7
  • Thermodynamics First Law of Thermodynamics Energy is neither created nor destroyed; it is transformed and/or transferred. DEuniverse = 0 Thermochemical equation includes heat changes which accompany a chemical reaction CH4(g) + 3 Cl2(g) CHCl3(g) + 3HCl(g) + 327 kJ heat given off 6-8
  • Figure 6.1 A chemical system and its surroundings. The system in this case is the contents of the reaction flask. The surroundings comprise everything else, including the flask itself. 6-9
  • The System and Its Surroundings A meaningful study of any transfer of energy requires that we first clearly define both the system and its surroundings. System + Surroundings = Universe The internal energy, E, of a system is the sum of the potential and kinetic energies of all the particles present. The total energy of the universe remains constant. A change in the energy of the system must be accompanied by an equal and opposite change in the energy of the surroundings. 6-10
  • Figure 6.2 Energy diagrams for the transfer of internal energy (E) between a system and its surroundings. Copyright The McGraw-Hill Companies, Inc. Permission required for reproduction or display. DE = Efinal - Einitial = Eproducts - Ereactants 6-11
  • Figure 6.3 The two cases where energy is transferred as heat only. Copyright The McGraw-Hill Companies, Inc. Permission required for reproduction or display. The system releases heat The system absorbs heat 6-12
  • Figure 6.4A The two cases where energy is transferred as work only. Zn(s) + 2H+(aq) + 2Cl-(aq) H2(g) + Zn2+(aq) + 2Cl-(aq) The system does work on the surroundings. 6-13
  • Figure 6.4B The two cases where energy is transferred as work only. The system has work done on it by the surroundings. 6-14
  • Table 6.1 The Sign Conventions* for q, w, and DE q + w = DE + + + + - depends on sizes of q and w - + depends on sizes of q and w - - - * For q: + means system gains heat; - means system releases heat. * For w: + means word done on system; - means work done by system. 6-15
  • The Law of Energy Conservation The first law of Thermodynamics states that the total energy of the universe is constant. Energy is conserved, and is neither created nor destroyed. Energy is transferred in the form of heat and/or work. DEuniverse = DEsystem + DEsurroundings = 0 6-16
  • Units of Energy The SI unit of energy is the joule (J). 1 J = 1 kgm2/s2 The calorie was once defined as the quantity of energy needed to raise the temperature of 1 g of water by 1 C. 1 cal = 4.184 J The British Thermal Unit (Btu) is often used to rate appliances. 1 Btu is equivalent to 1055 J. 6-17
  • Units of Energy Joule = 1 kgm2 / sec2 (Physics definition) calorie = heat to raise 1.00 g of water 1o C 1 calorie = 4.184 J (exact) 1 kilocalorie is what we use to measure the energy in our food. 6-18
  • Figure 6.5 Some quantities of energy. 6-19
  • Enthalpy: Chemical Change at Constant Pressure DE = q + w To determine DE, both heat and work must be measured. The most common chemical work is PV work the work done when the volume of a system changes in the presence of an external pressure. Enthalpy (H) is defined as E + PV so DH = DE + DPV If a system remains at constant pressure and its volume does not change much, then - DH DE 6-20
  • Figure 6.6 Two different paths for the energy change of a system. Even though q and w for the two paths are different, the total DE is the same for both. 6-21
  • Figure 6.7 Pressure-volume work. An expanding gas pushing back the atmosphere does PV work (w = -PDV). 6-22
  • H as a measure of E DH is the change in heat for a system at constant pressure. qP = DE + PDV = DH DH DE for reactions that do not involve gases for reactions in which the total amount (mol) of gas does not change for reactions in which qP is much larger than PDV, even if the total mol of gas does change. 6-23
  • Figure 6.8 Enthalpy diagrams for exothermic and endothermic processes. Copyright The McGraw-Hill Companies, Inc. Permission required for reproduction or display. CH4(g) + 2O2(g) CO2(g) + H2O(g) A 6-24 Exothermic process Heat is given out. H2O(s) H2O(l) B Endothermic process Heat is taken in.
  • Sample Problem 6.2 Drawing Enthalpy Diagrams and Determining the Sign of H PROBLEM: In each of the following cases, determine the sign of DH, state whether the reaction is exothermic or endothermic, and draw and enthalpy diagram. (a) H2(g) + O2(g) H2O(l) + 285.8 kJ (b) 40.7 kJ + H2O(l) H2O(g) PLAN: From each equation, note whether heat is a reactant or a product. If heat is taken in as a reactant, the process is endothermic. If heat is released as a product, the process is exothermic. For the enthalpy diagram, the arrow always points from reactants to products. For endothermic reactions, the products are at a higher energy than the reactants, since the reactants take in heat to form the products. 6-25
  • Sample Problem 6.2 SOLUTION: (a) H2(g) + O2(g) H2O(l) + 285.8 kJ Heat is a product for this reaction and is therefore given out, so the reaction is exothermic. The reactants are at a higher energy than the products. Energy H2(g) + O2(g) (reactants) 6-26 H = -285.8 kJ H2O(l) (products) EXOTHERMIC
  • Sample Problem 6.2 SOLUTION: (b) 40.7 kJ + H2O(l) H2O(g) Heat is a reactant in this reaction and is therefore absorbed, so the reaction is endothermic. The reactants are at a lower energy than the products. Energy H2O(g) 6-27 (reactants) H = + 40.7 kJ H2O(l) (products) ENDOTHERMIC
  • Enthalpy Changes Heat change DH > 0 (positive) ENDOTHERMIC Heat is absorbed by the process DH < 0 (negative) EXOTHERMIC Heat is produced by the process Competency VI-3 6-28
  • Calorimetry q = c x m x DT q = heat lost or gained c = specific heat capacity m = mass in g DT = Tfinal - Tinitial The specific heat capacity (c) of a substance is the quantity of heat required to change the temperature of 1 gram of the substance by 1C . 6-29
  • Heat Transfer of Liquids Specific Heat The amount of heat required to raise the temperature of one gram of a substance one degree Celsius(Kelvin), with no change in state. Water Ice Air Alcohol Metals 6-30 4.18 J/gC 2.03 J/gC 1.01 J/gC 2.46 J/gC 0.10 to 1.00 J/gC
  • Table 6.2 Specific Heat Capacities (c) of Some Elements, Compounds, and Materials Substance Specific Heat Capacity (J/gC) Elements Substance Specific Heat Capacity (J/gC) Solid materials aluminum, Al 0.900 wood 1.76 graphite,C 0.711 cement 0.88 iron, Fe 0.450 glass 0.84 copper, Cu 0.387 granite 0.79 gold, Au 0.129 steel 0.45 Compounds wat