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ENERGY (THERMOCHEMISTRY)______ Ch 1.5, 6, 9.10, 11.5-11.7, 12.3 Thermochemistry ̶ Prediction and measurement of energy transfer, in the form of heat, that accompanies chemical and physical processes. Chemical Thermodynamics (with Thermochemistry) Addresses the Questions 1. Will a particular reaction occur? 2. If it does occur then a) what energy changes and transfers are involved? b) to what extent? Energy (capacity to do work and/or cause heat transfer) kinetic (motion) potential (position, chemical composition) energy can be transferred from one form to another interconversion of potential and kinetic energy interconversion of work and heat heat, q (E transferred between two bodies at different T) work, w (E used to cause motion against an external force) 1st Law of Thermodynamics ̶ Conservation of Energy: Energy can neither be created nor destroyed, only converted from one form to another. internal energy, E ΔE = q + w ̶ heat adsorbed by system and work done on system FIG I ̶ Energy Transfer Between a System and its Surroundings -2- Definitions universe system surroundings system – characterized with respect to transfers of matter/energy with surroundings open closed isolated surroundings boundaries (walls) – characterized with respect to energy transfers in the form of heat adiabatic diathermal rigid/nonrigid WHY? ………………………………………. ability to predict REACTANTS initial state PRODUCTS final state prediction => state function => energy (ΔE?) => heat and work state function A State Function: Change in altitude depends only on the difference between initial and final values, not on the path. GOAL: if we can find the correct state function we can use it to predict the progress of a chemical reaction. Is Work the Desired State Function? (consider work in an ideal gas expansion) FIG II - Gas Confined to a Cylinder with a Movable Piston FIG III - Calculating the Work I ) I f 4 T = O -3- EX 1. For the combustion of ethanol at room temperature in the open air CH3CH2OH(l) + O2(g) → CO2(g) + H2O(l) What is the work done by the reaction? What is doing the work? What assumptions are you making? Is Heat the Desired State Function? (consider a beaker of hot or cold water) FIG IV - Heat Transfer Between a Beaker of Water and its Surroundings Working with Heat Transfer- Calorimetry: Measurement of quantities of heat heat at constant pressure, qP heat at constant volume, qV exothermic, endothermic adiabatic (Greek, adiabatos – impassable) thermal equilibrium heat capacity (rule of Dulong and Petit) -4- Heat Capacities of Some Substances (Table 6.4): cP = molar mass × cs Substance Specific Heat cs (J K-1 g-1) Hg(l) 0.1393 Al(s) Fe(s) Cu(s) 0.900 0.4493 0.3849 He(g) H(g) H2(g) 5.193 20.62 14.30 Molar Mass (g mol-1) 200.59 26.982 55.845 63.546 4.0026 1.0079 2.0158 N2(g) CH4(g) CaCO3(s) CH3COOH(l) 1.04 2.22 28.014 16.043 0.818 2.05 100.086 60.052 H2O(s) H2O(l) H2O(g) 2.1 4.184 2.0 18.015 18.015 18.015 Molar Heat Capacity cP ( J K-1 mol-1) 27.94 24.3 25.09 24.46 20.79 20.78 28.82 29.14 35.6 81.9 123.1 38 75.38 36 EX 2. 6.42 g of He absorbs 500.0 kJ of heat (at constant pressure) and the temperature increases by 15.0 K. What is CP? What is cs? What is cP? -5- Calculations with Heat Transfer calorimeter EX 3. 5.00 g of Fe (cs = 0.449 J K-1 g-1) at 400 K is put into 10.0 g of water (cs = 4.18 J K-1g-1) at 290 K. a) Ignoring the heat capacity of the calorimeter, what is the final temperature? b) If the calorimeter has a heat capacity of 500 J K-1 what is the final temperature? EX 4. A styrofoam cup calorimeter contains 150 g of water. Generation of 1430 J of heat inside the calori meter causes the temperature to increase by 1.93°C. What is Ccal? -6- Enthalpy (Greek, enthalpein – to warm in), H (heat at constant pressure, qp) enthalpy of reaction N2 (g) + 3 H2 (g) → 2 NH3(g) per 2 N2 (g) + 6 H2 (g) → 4 NH3 (g) 2 NH3 ΔH = -91. 8 kJ one mole of N2 three moles of H2 two moles of NH3 → N2 (g) + 3 H2(g) ΔH = -2(91.8) kJ ΔH = +91.8 kJ enthalpy of phase changes FIG V- Energy Changes Accompanying Phase Changes (Greek, phasis- appearance) Common phase transitions involve a transfer of heat between a system and its surroundings at constant temperature and pressure. 11.5 vaporization g → l (condensation l → g) 11.6 fusion (melting) s → l (freezing l → s) 11.7 sublimation s → g (deposition g → s) I Internal Energy, E (heat at constant volume, qV) Relation between ΔH and ΔE EX 5. At atmospheric pressure and 25°C the detonation of nitroglycerin 4 C 3 H 5 ( N O 3 ) 3 ( l) → 6 N 2 ( g ) + 1 0 H 2 O( g ) + 1 2 C O 2 ( g ) + O 2 ( g ) releases 5720 kJ of heat. What is the internal energy change for this reaction? -7- Experimental Calorimetry heat at constant pressure, qP FIG VI - Coffee Cup Calorimeter heat at constant volume, qV FIG VII - An Adiabatic Bomb Calorimeter EX 6. 0.1584 g of benzoic acid (C6H5COOH) are combusted in a constant volume bomb calorimeter. The temperature of the calorimeter rises by 2.54°C. ΔE for the combustion of benzoic acid is -26.38 kJ g-1. Determine the heat capacity of the calorimeter. Hess's Law - Extremely USEFUL combining enthalpies of reaction – enthalpy is a state function Hess’s Law: The change in enthalpy for a stepwise process is the sum of the changes in enthalpy of the steps. A + 2B → C C → 2D application to chemical reactions EX 7. Find the enthalpy change that accompanies the combustion of graphite to carbon monoxide from: ΔH = -393.5 kJ mol-1 1) C(s, gr) + O2(g) → CO 2 (g) ΔH = 283.0 kJ mol-1 2) CO 2 (g) → CO(g) + 1/2 O 2 (g) C(s, gr) + 1/2 O 2 (g) → CO(g) ΔH = ? -8- FIG VIII- Enthalpy Diagrams for Chemical Reactions application to phase changes FIG IX -Enthalpy Change upon Cooling 2.5 mol of Water from 130°C to -40°C EX 8. How much heat is required to lower the temperature of 2.5 mol of water from 130°C to -40°C? For water: cP(s), cP(l), cP(g) = 37.6, 75.38, 33.1 J mol-1 K-1 , ΔHvap, ΔHfus = 40.7, 6.02 kJ mol-l overall: H2O(g, 130°C) → H2O(s, -40°C): a) H2O(g, 130°C) → H2O(g, 100°C) : b) H2O(g, 100°C) → H2O(l, 100°C) : c) H2O(l, 100°C) → H2O(l, °C) : d) H2O(l, °C) → H2O(s, -0°C) : e) H2O(s, 0°C) → H2O(s, -40°C) : -9- "Standard" Enthalpies standard states – stable form (allotrope) at P = 1 atm and specified temperature (usually T = 25°C) pure solid pure liquid gas – ideal gas behavior 1 molar (1 M) solution – ideal solution behavior standard enthalpy of formation, ∆Hfo, of a compound: enthalpy change for the reaction forming one mole of the compound from its elements in their standard states. ∆Hfo (element in standard state) = 0 EX 9. Find the enthalpy change that accompanies the combustion of methane to carbon dioxide from enthalpies of formation (∆Hfo in Appendix IIB) defining formation reaction ∆Hfo (kJ mol-1) -74.81 1) C(s, gr) + 2 H2(g) → CH4(g) O2(g) → O2(g) 3) C(s, gr) + O2(g) → CO2(g) -393.51 4) H2(g) + 1/2 O2(g) → H2O (l) -285.83 2) 0 TARGET: CH4(g) + 2 O2(g) → CO2(g) + 2 H2O(l) long way: short way: FOR ANY REACTION: ∆Horx = ∑.nprod ∆Hof (products) – ∑.nreact ∆Hof (reactants) => products - reactants - 10 - application to a) chemical reactions and b) phase changes EX 10. For the combustion of ethanol (unbalanced) CH 3 CH 2 O H( l) + O 2 ( g ) → CO 2 ( g ) + H 2 O( l ) a) Determine the standard enthalpy of the reaction given that ∆Hfo for CH 3CH 2OH(l), CO 2(g), and H 2O(l) are -277.0, -393.509, and -285.83, respectively. b) What is the standard enthalpy change when all reactants and products are gases? At 25°C, ΔHvap(H 2O) = 44.01 kJ mol-1 and ΔHvap (CH 3CH 2OH) = 42.59 kJ mol-1. application to combustion reaction/ignoring the work term EX 11. Determine ∆Hof of benzoic acid (C 6 H 5 C OO H). ΔE o for the combustion of benzoic acid is -26.38 kJ g-1,∆Hof [CO 2(g)] = -393.5 kJ mol-1 ∆Hof [H 2O(1)] = -285.840 kJ mol-1. - 11 - Average Bond Enthalpies (Chapter 9.10- Bond Energies) Heat, at standard conditions, absorbed by a system to break chemical bonds in the gas phase. Bond Energy of a Homonuclear Diatomic bond enthalpy of chlorine gas (measured, 243 kJ/mol): Cl2(g) ½ Cl2(g) → 2 Cl(g) → Cl(g) ΔHo = ? ΔHo = 121.68 2 × 121.68 = 243.46, close to 243 Bond Energy of Methane 1) atomization enthalpy of CH4 (energy to break all bonds and form free atoms, atomize, in gas phase): CH4(g) → ΔHo = 1663 kJ/mol per CH, 1663/4 = 415 C(g) + 4 H(g) 2) enthalpy of a C – H bond in CH4 (experimental results): CH4(g) CH3(g) CH2(g) CH(g) → CH3(g) + H(g) → CH2(g) + H(g) → CH(g) + H(g) → C(g) + H(g) ΔHo = 439 kJ/mol ΔHo = 453 ΔHo = 425 ΔHo = 339 average = 1656/4 = 414 3) enthalpy of a C – H bond: CH4(g) C2H6(g) CHF3(g) CHBr3(g) → → → → ΔHo = 439 kJ/mol ΔHo = 410 ΔHo = 429 ΔHo = 377 above 4 within 8% of average CH3(g) + H(g) C2H5(g) + H(g) CF3(g) + H(g) CBr3(g) + H(g) Table 9-3. Average Bond Enthalpies Single Bonds H C N O F Cl Br I S Si H 436 412 388 463 565 431 366 299 338 318 C 348 305 360 484 338 276 238 259 N O F 163 157 270 200 146 185 203 158 254 496 Cl Br 243 219 210 250 193 178 212 I Si S 151 264 226 466 Multiple Bonds C=C C≡C O2 612 812 497 N=N N≡N CO 2 409 946 799 C=N C≡N 613 890 C=O C≡O 743 1046 - 12 - FIG X. Estimate ΔHo for: 2) 1) -676 4) -824 5) -1130 6) 872 968 PRODUCTS REACTANTS 3) CC12F2 + 2 H 2 → CH 2Cl 2 + 2 HF 676 2516 -2630 => reactants - products EX 12. Use average bond enthalpies to estimate ΔHo for 2 CH 2 = CH – CH 3 (g) + 2 NH 3 (g) + 3 O 2 (g) → 2 CH 2 = CH – CN(g) + 6 H 2O(g)