THERMOCHEMISTRY Thermodynamics The study of Heat and Work and State Functions

1. THERMOCHEMISTRYThermodynamicsThe study of Heat and Work and State Functions

2. 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

3. Burning peanuts supply sufficient energy to boil a cup of water. Burning sugar (sugar reacts with KClO3, a strong oxidizing agent) Energy & Chemistry

4. These reactions are PRODUCT FAVORED They proceed almost completely from reactants to products, perhaps with some outside assistance. Energy & Chemistry

5. 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- CCR, page 845

6. Potential & Kinetic Energy Potential energy — energy a motionless body has by virtue of its position.

7. Potential Energyon 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.

8. Potential Energyon 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

9. Potential & Kinetic Energy Kinetic energy — energy of motion • Translation

10. Potential & Kinetic Energy Kinetic energy — energy of motion.

11. 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

12. Internal Energy (E) • PE + KE = Internal energy (E or U)

13. Internal Energy (E) • The higher the T the higher the internal energy • So, use changes in T (∆T) to monitor changes in E (∆E).

14. Thermodynamics • Thermodynamics is the science of heat (energy) transfer. Heat energy is associated with molecular motions. Heat transfers until thermal equilibrium is established.

15. T(system) goes down T(surr) goes up Directionality of Heat Transfer • Heat always transfer from hotter object to cooler one. • EXOthermic: heat transfers from SYSTEM to SURROUNDINGS.

16. T(system) goes up T (surr) goes down Directionality of Heat Transfer • Heat always transfer from hotter object to cooler one. • ENDOthermic: heat transfers from SURROUNDINGS to the SYSTEM.

17. Energy & Chemistry All of thermodynamics depends on the law of CONSERVATION OF ENERGY. • The total energy is unchanged in a chemical reaction. • If PE of products is less than reactants, the difference must be released as KE.

18. Energy Change in Chemical Processes PE of system dropped. KE increased. Therefore, you often feel a T increase.

19. James Joule 1818-1889 UNITS OF ENERGY 1 calorie = heat required to raise temp. of 1.00 g of H2O by 1.0 oC. 1000 cal = 1 kilocalorie = 1 kcal 1 kcal = 1 Calorie (a food “calorie”) But we use the unit called the JOULE 1 cal = 4.184 joules

20. Which has the larger heat capacity? HEAT CAPACITY The heat required to raise an object’s T by 1 ˚C.

21. Specific Heat Capacity How much energy is transferred due to T difference? The heat (q) “lost” or “gained” is related to a) sample mass b) change in T and c) specific heat capacity

22. Specific Heat Capacity Substance Spec. Heat (J/g•K) H2O 4.184 Ethylene glycol 2.39 Al 0.897 glass 0.84 Aluminum

23. Specific Heat Capacity If 25.0 g of Al cool from 310 oC to 37 oC, how many joules of heat energy are lost by the Al?

24. heat gain/lose = q = (sp. ht.)(mass)(∆T) Specific Heat Capacity If 25.0 g of Al cool from 310 oC to 37 oC, how many joules of heat energy are lost by the Al? where ∆T = Tfinal - Tinitial q = (0.897 J/g•K)(25.0 g)(37 - 310)K q = - 6120 J Notice that the negative sign on q signals heat “lost by” or transferred OUT of Al.

25. Heat TransferNo Change in State q transferred = (sp. ht.)(mass)(∆T)

26. Heat Transfer with Change of State Changes of state involve energy (at constant T) Ice + 333 J/g (heat of fusion) -----> Liquid water q = (heat of fusion)(mass)

27. Heat Transfer and Changes of State Requires energy (heat). This is the reason a) you cool down after swimming • you use water to put out a fire. Liquid ---> Vapor + energy

28. Heating/Cooling Curve for Water Evaporate water Heat water Note that T is constant as ice melts Melt ice

29. +333 J/g +2260 J/g Heat & Changes of State What quantity of heat is required to melt 500. g of ice and heat the water to steam at 100 oC? Heat of fusion of ice = 333 J/g Specific heat of water = 4.2 J/g•K Heat of vaporization = 2260 J/g

30. Heat & Changes of State How much heat is required to melt 500. g of ice and heat the water to steam at 100 oC? 1. To melt ice q = (500. g)(333 J/g) = 1.67 x 105 J 2. To raise water from 0 oC to 100 oC q = (500. g)(4.2 J/g•K)(100 - 0)K = 2.1 x 105 J 3. To evaporate water at 100 oC q = (500. g)(2260 J/g) = 1.13 x 106 J 4. Total heat energy = 1.51 x 106 J = 1510 kJ

31. Chemical Reactivity What drives chemical reactions? How do they occur? The first is answered by THERMODYNAMICS and the second by KINETICS. Have already seen a number of “driving forces” for reactions that are PRODUCT-FAVORED. • formation of a precipitate • gas formation • H2O formation (acid-base reaction) • electron transfer in a battery

32. Chemical Reactivity But energy transfer also allows us to predict reactivity. In general, reactions that transfer energy to their surroundings are product-favored. So, let us consider heat transfer in chemical processes.

33. Heat Energy Transfer in a Physical Process CO2 (s, -78 oC) ---> CO2 (g, -78 oC) Heat transfers from surroundings to system in endothermic process.

34. Heat Energy Transfer in a Physical Process • CO2 (s, -78 oC) ---> CO2 (g, -78 oC) • A regular array of molecules in a solid -----> gas phase molecules. • Gas molecules have higher kinetic energy.

35. CO2 gas ∆E = E(final) - E(initial) = E(gas) - E(solid) CO2 solid Energy Level Diagram for Heat Energy Transfer

36. Heat Energy Transfer in Physical Change • Gas molecules have higher kinetic energy. • Also, WORKis done by the system in pushing aside the atmosphere. CO2 (s, -78 oC) ---> CO2 (g, -78 oC) Two things have happened!

37. heat energy transferred work done by the system energy change FIRST LAW OF THERMODYNAMICS ∆E = q + w Energy is conserved!

38. heat transfer in (endothermic), +q heat transfer out (exothermic), -q w transfer in (+w) w transfer out (-w) SYSTEM ∆E = q + w

39. ENTHALPY Most chemical reactions occur at constant P, so Heat transferred at constant P = qp qp = ∆H where H = enthalpy and so ∆E = ∆H + w (and w is usually small) ∆H = heat transferred at constant P ≈ ∆E ∆H = change in heat content of the system ∆H = Hfinal - Hinitial

40. ENTHALPY ∆H = Hfinal - Hinitial If Hfinal > Hinitial then ∆H is positive Process is ENDOTHERMIC If Hfinal < Hinitial then ∆H is negative Process is EXOTHERMIC

41. USING ENTHALPY Consider the formation of water H2(g) + 1/2 O2(g) --> H2O(g) + 241.8 kJ Exothermic reaction — heat is a “product” and ∆H = – 241.8 kJ

42. USING ENTHALPY Making liquid H2O from H2 + O2 involves twoexothermic steps. H2 + O2 gas H2O vapor Liquid H2O

43. USING ENTHALPY Making H2O from H2 involves two steps. H2(g) + 1/2 O2(g) ---> H2O(g) + 242 kJ H2O(g) ---> H2O(liq) + 44 kJ ----------------------------------------------------------------------- H2(g) + 1/2 O2(g) --> H2O(liq) + 286 kJ Example of HESS’S LAW— If a rxn. is the sum of 2 or more others, the net ∆H is the sum of the ∆H’s of the other rxns.

44. Hess’s Law & Energy Level Diagrams Forming H2O can occur in a single step or in a two steps. ∆Htotal is the same no matter which path is followed.

45. Hess’s Law & Energy Level Diagrams Forming CO2 can occur in a single step or in a two steps. ∆Htotal is the same no matter which path is followed.

46. ∆H along one path = ∆H along another path • This equation is valid because ∆H is a STATE FUNCTION • These depend only on the state of the system and not on how the system got there. • V, T, P, energy — and your bank account! • Unlike V, T, and P, one cannot measure absolute H. Can only measure ∆H.

47. Standard Enthalpy Values Most ∆H values are labeled ∆Ho Measured under standard conditions P = 1 bar = 105 Pa = 1 atm /1.01325 Concentration = 1 mol/L T = usually 25 oC with all species in standard states e.g., C = graphite and O2 = gas

48. Enthalpy Values Depend on how the reaction is written and on phases of reactants and products H2(g) + 1/2 O2(g) --> H2O(g) ∆H˚ = -242 kJ 2 H2(g) + O2(g) --> 2 H2O(g) ∆H˚ = -484 kJ H2O(g) ---> H2(g) + 1/2 O2(g) ∆H˚ = +242 kJ H2(g) + 1/2 O2(g) --> H2O(liquid) ∆H˚ = -286 kJ

49. Standard Enthalpy Values NIST (Nat’l Institute for Standards and Technology) gives values of ∆Hfo = standard molar enthalpy of formation — the enthalpy change when 1 mol of compound is formed from elements under standard conditions. See Table 6.2

50. ∆Hfo, standard molar enthalpy of formation Enthalpy change when 1 mol of compound is formed from the corresponding elements under standard conditions H2(g) + 1/2 O2(g) --> H2O(g) ∆Hfo (H2O, g)= -241.8 kJ/mol By definition, ∆Hfo= 0 for elements in their standard states.