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Chapter 6: Energy and Chemical Reactions

Chapter 6: Energy and Chemical Reactions. Energy Types of energy Energy units Specific heat and energy transfer Changes of state Enthalpy: Thermochemical Equations Definitions Enthalpies of reaction Hess's law Enthalpy of formation standard state; standard molar enthalpies of formation

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Chapter 6: Energy and Chemical Reactions

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  1. Chapter 6: Energy and Chemical Reactions

  2. Energy • Types of energy • Energy units • Specific heat and energy transfer • Changes of state • Enthalpy: Thermochemical Equations • Definitions • Enthalpies of reaction • Hess's law • Enthalpy of formation • standard state; standard molar enthalpies of formation • enthalpies of reaction from enthalpies of formation • Calorimetry • Coffee-cup calorimetry • Bomb calorimetry

  3. Energy = capacity to do work or transfer heat chemical energy in gasoline & O2 move car & generate heat (w = F•d)  A. Types of energy • Kinetic energy: energy of motion • mechanical: motion of objects (KE = ½mv2) • thermal: motion of atoms and molecules • electrical: motion of electrons • Potential energy: stored energy • gravitational energy (PE = mgh) • chemical energy: energy stored in chemical bonds • electrostatic energy: attraction or repulsion of charged particles  Total energy = KE + PE conserved! (law of conservation of energy)

  4. Energy • Energy units • SI: define 1 joule (J) = KE of a 2-kg object traveling at 1 m/s • KE = ½mv2 = ½(2 kg)(1 m/s)2 = 1 kg·m2/s2 = 1 J • 1 kJ = 1000 J • English: 1 cal = heat required to raise the temperature of 1 g of water by 1ºC • 1 kcal = 1000 cal • 1 cal = 4.184 J • 1 kcal = 4.184 kJ 1 Calorie (Cal) = 1000 cal = 1 kcal (food calorie)

  5. cal g ·ºC q m•T J g·K • Energy • Specific heat • heat capacity: heat required to raise the temperature of something • (extensive) by 1ºC (1 K) • specific heat: heat required to raise the temperature of 1 g of something • (intensive) by 1ºC • for water, • specific heat, C = = or C = 1 cal/g ºC • or 4.184 J/g ºC • or rearranging for heat: q = C m  DT • How many kJ of heat are required to raise the temperature of 5.00 x 102 g of water by 24.0 ºC?

  6. Energy • Changes of state • fusion (melting): • solid + heat  liquid • qfusion = heat required to melt a substance • vaporization (boiling): • liquid + heat  gas • qvaporization = heat required to vaporize a substance

  7. Energy • Changes of state Heating/cooling curves: gas (e) liquid + gas b.p. (d) T liquid (c) solid + liquid m.p. (b) solid (a) heat added  (heating)  heat removed (cooling) e.g. H2O (a) C(s) = 2.1 J/g·ºC (d) qvap = 2256 J/g or 40.7 kJ/mol (b) qfus = 333.5 J/g or 6.01 kJ/mol (e) C(g) = 2.0 J /g·ºC (c) C(l) = 4.2 J /g·ºC

  8. Energy • Changes of state How much heat is required to convert 10.0 g of solid water at -10ºC to gas at 100 ºC?

  9. Enthalpy: Thermochemical Equations • Definitions universe: surroundings system (energy can be transferred) Internal energy, heat, and work: For a system: DEsys = qsys + wsys change in energy for a system heat absorbed or released by the system work done on or by the system

  10. Enthalpy: Thermochemical Equations • Definitions DEsys = qsys + wsys • Esys = energy of a system • = sum of all energies (kinetic, bond, etc.) • DEsys = Efinal - Einitial • system gains energy: DEsys > 0 (DEsurr < 0) • system loses energy: DEsys < 0 (DEsurr > 0) • qsys = heat • heat added to system: qsys > 0 endothermic (qsurr < 0) • heat lost from system: qsys < 0 exothermic (qsurr > 0) • wsys = work • work done on system: wsys > 0 (wsurr < 0) • work done by system: wsys < 0 (wsurr > 0) can’t measure can measure

  11. Enthalpy: Thermochemical Equations • Definitions A balloon is heated by adding 240 J of heat. It expands, doing 135 J of work on the atmosphere. What is DE for the system?

  12. 1 g H2O, 65 ºC 1 g H2O, 50 ºC DE = 25 cal 1 g H2O, 25 ºC 1 g H2O, 10 ºC • Enthalpy: Thermochemical Equations • Definitions • DE is a state function • -depends only on the initial and final states of a system, not the path • taken to get from one to the other First Law of Thermodynamics: energy is neither created nor destroyed; it can only change forms (conservation of energy). For a system: DEsys = qsys + wsys

  13. Enthalpy: Thermochemical Equations • Enthalpies of reaction • Enthalpy change, DH = heat gained or lost at constant pressure • DH = Hfinal - Hinitial = qP (does not include work) • -state function Enthalpy of reaction • Reactants  Products • DH = Hproducts - Hreactants • DH > 0: endothermic (products higher energy than reactants) • DH < 0: exothermic (products lower energy than reactants)

  14. Enthalpy: Thermochemical Equations • Enthalpies of reaction • e.g., 2Al + Fe2O3 2Fe + Al2O3DH = -851.5 kJ (thermochemical equation) • 1. DH is an extensive property • What is DH per mole of Fe? • 2. reverse equation: change sign of DH • 2Fe + Al2O3  2Al + Fe2O3DH = ? • 3. enthalpy diagrams: 2Al + Fe2O3 DH = –851.5 kJ DH = +851.5 kJ H 2Fe + Al2O3

  15. Enthalpy: Thermochemical Equations • Hess’s law DHoverall process = SDHindividual steps (by any path; 1st law) • 3Mg + N2 Mg3N2DH = ? • given: Mg3N2 + 3H2  3Mg + 2NH3DH = +371 kJ • and: 1/2N2 + 3/2H2  NH3DH = –46 kJ

  16. Enthalpy: Thermochemical Equations • Hess’s law • S + O2 SO2DH = ? • given: 2SO2 + O2  2SO3DH = –196 kJ • and: 2S + 3O2  2SO3DH = –790 kJ

  17. Enthalpy: Thermochemical Equations • Enthalpy of formation • 1. standard state = most stable form of a substance at 1 bar pressure • and a given temperature, usually 25ºC (298 K). • noble gases: monatomic gases • diatomic gases: brinclhof • metals: Mº • carbon: graphite • standard molar enthalpy of formation, DHfº • = enthalpy of reaction for formation of one mole of a substance • from its elements, all in their standard states (DHfº(element) = 0) • (Table 6.2, Appendix L) • e.g., DHfº(CO(g)) = –110.5 kJ/mol • means: C(graphite) + ½O2(g)  CO(g) DHrxnº = –110.5 kJ/mol • = DHfº

  18. Enthalpy: Thermochemical Equations • Enthalpy of formation • enthalpies of reaction from enthalpies of formation DHrxnº =  nDHfº(products) –  nDHfº(reactants) C2H2(g) + 2H2(g)  C2H6(g) DHrxnº = ? C2H5OH(l) + 3O2(g)  2CO2(g) + 3H2O(l) DHrxnº = ? DHfº (kJ/mol) C2H2(g) 226.7 C2H6(g) -84.68 C2H5OH(l) -277.7 CO2(g) -393.5 H2O(l) -285.8

  19. Enthalpy: Thermochemical Equations • Enthalpy of formation • enthalpies of reaction from enthalpies of formation • The heat of combustion of glucose is –2802.5 kJ. What is DHfº for glucose? • C6H12O6(s) +

  20. Calorimetry • Coffee-cup calorimetry thermometer stopper or lid • two styrofoam cups (insulation) • -enclose the universe: system (reaction) • and surroundings (solution) reaction in solution generates or consumes heat, raising or lowering temperature • qsolution = C m  T • Hrxn = qrxn = –qsolution • i.e., if T increases (T>0), rxn is exothermic (H<0) • if T decreases (T<0), rxn is endothermic (H>0)

  21. Calorimetry • Coffee-cup calorimetry When 4.25 g of NH4NO3(s) dissolves in 60.0 g of H2O in a coffee cup calorimeter, the temperatures drops from 22.0ºC to 16.9ºC. Assuming the specific heat of the solution is 4.184 J/g·ºC, what is H (in kJ/mol) for the dissolution of NH4NO3? (fw NH4NO3 80.04)

  22. Calorimetry • Bomb calorimetry thermometer heating element stirrer insulation heat capacity of calorimeter = C qrxn = -qcal = –C T bomb

  23. Calorimetry • Bomb calorimetry A 1.80-g sample of octane, C8H18, was combusted in a bomb calorimeter with a heat capacity of 11.66 kJ/ºC. The temperature increased from 21.36ºC to 28.78ºC. What is H of combustion per mole of octane? (mw C8H18 114.23)

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