Ch. 6 & 8.8 lecture notes

1. Ch. 6 & 8.8 lecture notes THERMOCHEMISTRY!!!

2. Video: Exploding Gummy Bear! I’m a wimp, so you’re not getting the live version!  Please don’t try this at home! 

3. Terms to know • System: • Surroundings: • Heat: • Work: • Energy:

4. Rvw: Endothermic vs. Exothermic Reactions

5. 1st Law of Thermodynamics • Nrg of universe is constant • Nrg is conserved; cannot be created nor destroyed • Internal nrg (E) = total nrg of the system. • can be changed by flow of work (w), heat (q), or both: ΔE = q + w • q > 0 = • q < 0 = • W > 0 = • W < 0 =

6. State functions • Internal nrg (E) is a state function • State functions depend only on initial and final state of the system • Independent of path taken to get from start to finish. • Example:

7. P-V work

8. P-V work • Work in terms of PV: • For expanding gas: • For compressed gas:

9. Enthalpy (H) • Relates nrg & P-V work; H = E + PV • At constant pressure (where only P-V work is allowed), ΔH = nrg flow as heat = qp • How is this so??? Derive the equation!

10. Enthalpy of rxns (ΔH ) • ΔH = qp = Hfinal – Hinitial = Hprods– Hrxts • Characteristics of enthalpy changes (ΔH ) • Enthalpy is an extensive property • ΔHforward = -ΔH reverse • ΔH depends on state of rxts and prods • In terms of potential nrg, …. • Assume everything is at standard conditions (25 oC, 1 atm) unless stated otherwise

11. Practice problem #1 2H2O2 (l)  2H2O (g) + O2 (g) ΔHrxn = -196 kJ What is ΔH if 5.00 g of H2O2 is decomposed?

12. Calorimetry • Measures heat flow by measuring . . . • Uses a calorimeter (see picture) • Constant-pressure calorimetry: • Patm remains constant • Used to determine changes in enthalpy (heats of rxn) for rxnsoccuring in solution •  ΔH = qp 

13. Coffee Cup Calorimeter On the outside… On the inside…

14. Constant-volume calorimetry: A “bomb” calorimeter • No work is done. Why? • Weighed rxts placed inside rigid steel container (the “bomb”) and ignited • ΔH = q + w = qV

15. Heat capacity • Heat capacity (C): nrg needed to raise temp of a body by….. Units = ???? • C = heat absorbed • increase in temp • Molar heat capacity: heat capacity of one of substance; Units = ???? • Specific heat capacity (c): heat capacity of one of substance; Units = ????

16. Some specific heats Substance          Specific Heat at 25oC   (J/goC) Lead                           0.128Gold                            0.129 Silver                          0.235Copper                       0.387 Iron                             0.4498 Carbon (graphite)        0.711Granite                        0.803 Olive Oil                     2.0Ethyl alcohol                2.45Water, (liquid)            4.1796 Notice anything about the first 5 substances?

17. Heat transfer

18. Heat transfer • Q = mcΔT = ΔH if at constant P • If a hot object touches a cooler object… • TfTi for hot object • TfTi for cooler object • And Tf of hot object Tf of cooler object • All heat lost by one object is gained by the other, so… • Q1= • Tip to avoid dealing with negative…

19. Practice problem #2 • A 150.0 g sample of a metal at 75.0 oC is added to 150.0 g of water at 15.0 oC. The temperature of the water rises to 18.3 oC. Calculate the specific heat of the metal, assuming that all of the heat lost by the metal is gained by the water. (specific heat of water = 4.18 J/g oC)

20. Phase changes & enthalpy

21. Phase changes & enthalpy • How do you calculate enthalpy for segments BC and DE? • How do you calculate enthalpy for segments AB, CD and EF?

22. Phase change constants for water • Specific heat (c) for ice = 2.09 J/g•oC • Specific heat (c) for water = 4.18 J/g•oC • Specific heat (c) for vapor = 1.84 J/g•oC • ΔHfus = Heat of fusion = 6.01 kJ/mol • ΔHvap = Heat of vaporization = 40.67 kJ/mol

23. Phase change calorimetry

24. Practice Problems • Do problems on Calorimetry wkhst

25. Pre-assessment: Lewis Structures • On your white board, draw the Lewis structures for the following molecules: • H2O • CO2 • HCN

26. Bond nrg & enthalpy of rxns • Reactants break bonds and new bonds form to make products • Breaking bonds is…. • Forming bonds is…. • Bond enthalpy(D): • As bond strength increases, bond enthalpy…. • Use bond enthalpies to calculate ΔH without having to know ΔHf’s for all species

27. More on bond enthalpies • Need to know Lewis structures for compounds. Just seeing “N2H4” doesn’t help us calculate the bond enthalpy for it. We need to know HOW it bonds: So, there are N-H bonds and N-N bond

28. ΔHrxn = Σ (D of bonds broken) - Σ (D of bonds formed) • NOTE: Sum of RXTS – Sum of PRODUCTS!!! • Multiply bond enthalpy by number of bonds • If ΔHrxn> 0, nrg to break bonds is nrg released when new bonds form • Vice-versa if ΔHrxn < 0.

29. Bond enthalpy, bond length and bond strength

30. Interpreting the chart • As bond length , bond enthalpy • As bond length , # of bonds • As # of bonds , bond enthalpy • As bond strength , bond enthalpy • Therefore: • As bond length , bond strength • As # of bonds , bond strength

31. Bond enthalpies for practice probs BondBond enthalpy (kJ/mol) C—C 347 C=C 614 C C 839 H—H 432 N — N 160 N — H 391

32. BondBond enthalpy (kJ/mol) N  N 941 Cl — Cl 239 C — H 413 H — Cl 427 C — Cl 339

33. Practice problem #3 Calculate ΔHrxn for each of the following: • Cl2 + CH4 CH3Cl + HCl • N2H4 N2 + 2 H2 • C2H4 + HCl  C2H3Cl + H2

34. Practice Problem #4 • The standard enthalpy of formation of NH3 is -46 kJ/mol. Use this information as well as the balanced equation below to estimate the N-H bond energy. Compare your result with the value in previous table. N2 (g) + 3H2 (g) → 2NH3 (g)