1. FINAL EXAM Wednesday,December 11, at 10:15 a.m. – 12:15 p.m. in the IC building, Room 421

2. Prentice Hall © 2003Chapter 11 The forces holding solids and liquids together are called intermolecular forces. The covalent bond holding a molecule together is an intramolecular forces. The attraction between molecules is an intermolecular force. Intermolecular forces are much weaker than intramolecular forces When a substance melts or boils the intermolecular forces are broken (not the covalent bonds). Intermolecular Forces

3. Prentice Hall © 2003Chapter 11 Intermolecular Forces

4. Prentice Hall © 2003Chapter 11 Ion-Dipole Forces Interaction between an ion and a dipole (e.g. water). Strongest of all intermolecular forces. Intermolecular Forces

5. Prentice Hall © 2003Chapter 11 Dipole-Dipole Forces Exist between neutral polar molecules. Polar molecules need to be close together. Weaker than ion-dipole forces. There is a mix of attractive and repulsive dipole-dipole forces as the molecules tumble. If two molecules have about the same mass and size, then dipole-dipole forces increase with increasing polarity. Intermolecular Forces

6. Prentice Hall © 2003Chapter 11 Dipole-Dipole Forces Intermolecular Forces

7. Prentice Hall © 2003Chapter 11 Dipole-Dipole Forces Intermolecular Forces

8. Prentice Hall © 2003Chapter 11 London Dispersion Forces Weakest of all intermolecular forces. It is possible for two adjacent neutral molecules to affect each other. The nucleus of one molecule (or atom) attracts the electrons of the adjacent molecule (or atom). For an instant, the electron clouds become distorted. In that instant a dipole is formed (called an instantaneous dipole). Intermolecular Forces

9. Prentice Hall © 2003Chapter 11 London Dispersion Forces Polarizability is the ease with which an electron cloud can be deformed. The larger the molecule (the greater the number of electrons) the more polarizable. London dispersion forces increase as molecular weight increases. London dispersion forces exist between all molecules. London dispersion forces depend on the shape of the molecule. Intermolecular Forces

10. Prentice Hall © 2003Chapter 11 London Dispersion Forces The greater the surface area available for contact, the greater the dispersion forces. London dispersion forces between spherical molecules are lower than between sausage-like molecules. Intermolecular Forces

11. Prentice Hall © 2003Chapter 11 London Dispersion Forces Intermolecular Forces

12. Prentice Hall © 2003Chapter 11 London Dispersion Forces Intermolecular Forces

13. Prentice Hall © 2003Chapter 11 Hydrogen Bonding Special case of dipole-dipole forces. By experiments: boiling points of compounds with H-F, H-O, and H-N bonds are abnormally high. Intermolecular forces are abnormally strong. Intermolecular Forces

14. Prentice Hall © 2003Chapter 11 Hydrogen Bonding

15. Prentice Hall © 2003Chapter 11 Intermolecular Forces

16. Sublimation: solid  gas. Vaporization: liquid  gas. Melting or fusion: solid  liquid. Deposition: gas  solid. Condensation: gas  liquid. Freezing: liquid  solid. Phase Changes

19. Energy Changes Accompanying Phase Changes  H sub > 0 (endothermic).  H dep < 0 (exothermic). Phase Changes

20. . ENERGY ASSOCIATED WITH HEATING CURVES

21. Topics Vapor Pressure Normal Boiling Point Normal Freezing Specific Heat Enthalpy (Heat) of Vaporization Enthalpy (Heat) of Fusion

22. Vapor Pressure THE PRESSURE OF A VAPOR IN EQUILIBRIUM WITH ITS LIQUID (OR ITS SOLID)

23. NORMAL BOILING POINT & FREEZING POINTS NORMAL BOILING PT. - THE TEMPERATURE @WHICH VAPOR PRESSURE = 1 atm NORMAL FREEZING PT. – THE TEMPERATURE @ WHICH THE VAPOR PRESSURE OF THE SOLID AND THE LIQUID ARE THE SAME

25. Heat Capacity aka Specific Heat (C) Specific Heat (C) = the amount of energy required to raise the temperature of 1 gram of substance 1 degree celcius

26. Specific Heat (C) aka Heat Capacity Units for: specific heat (C) = J/g- o C where J = joules o C = temperature in o C g = mass in grams

27. Specific Heat (C) Values (aka Heat Capacity) Example: Water LIQUID: C Liq = 4.18 J/ ( o C. g) LIQUID: C sol = 2.09 J/ ( o C. g) LIQUID: C gas = 1.84 J/ ( o C. g)

28. Use of Specific Heat q = mC  T q = gm substance x specific heat x  T where: M = mass of substance in grams q = amount of heat (energy) C = specific heat And  T = change in temperature

29. Enthalpy of Vaporization aka heat of vaporization (  H vap ) Is the amount of heat needed to convert a liquid to a vapor at its normal boiling point

30. Enthalpy of Fusion aka heat of fusion (  H fus ) Is the amount of heat needed to convert a solid to a liquid at its normal melting (freezing) point

31. Units for  H vap,  H fus and heat(q) Heat of fusion  H fus = kJ/mol Heat of vaporization  H vap = kJ/mol Heat (q) = Joules

32. Therefore: To come up with Joules which is the unit of heat, if:  H is given, then: q vap =  H vap x moles and q fus =  H fus x moles (2)Specific heat (C) is given, then: q = mC  T

34. Sample Problem Calculate the enthalpy change upon converting 1 mole of ice at -25 o C to steam at 125 o C under a constant pressure of 1 atm? The specific heats are of ice, water and steam 2.09 J/g-K for ice, 4.18 J/g-K for water and 1.84 J/g-K for steam. For water,  H fus = 6.01 kJ/mol, and  H vap = 40.67kJ/mol. Note: The total enthalpy change is the sum of the changes of the individual steps.

35. Energy Changes Accompanying Phase Changes All phase changes are possible under the right conditions. The sequence heat solid  melt  heat liquid  boil  heat gas is endothermic. The sequence cool gas  condense  cool liquid  freeze  cool solid is exothermic. Phase Changes

36. Heating Curves Plot of temperature change versus heat added is a heating curve. During a phase change, adding heat causes no temperature change. Supercooling: When a liquid is cooled below its melting point and it still remains a liquid. Achieved by keeping the temperature low and increasing kinetic energy to break intermolecular forces. Phase Changes

37. HEATING CURVES ENERGY ASSOCIATED WITH HEATING CURVES During a phase change, adding heat causes no temperature change.

39. Critical Temperature and Pressure Gases liquefied by increasing pressure at some temperature. Critical temperature: the minimum temperature for liquefaction of a gas using pressure. Critical pressure: pressure required for liquefaction. Phase Changes

40. Phase diagram: plot of pressure vs. Temperature summarizing all equilibria between phases. Given a temperature and pressure, phase diagrams tell us which phase will exist. Any temperature and pressure combination not on a curve represents a single phase. Phase Diagrams

41. Features of a phase diagram: –Triple point: temperature and pressure at which all three phases are in equilibrium. –Vapor-pressure curve: generally as pressure increases, temperature increases. –Critical point: critical temperature and pressure for the gas. –Melting point curve: as pressure increases, the solid phase is favored if the solid is more dense than the liquid. –Normal melting point: melting point at 1 atm. Phase Diagrams

43. The Phase Diagrams of H 2 O and CO 2 Phase Diagrams

44. 3 Things Learned Reading a phase diagram Determining triple point on phase diagram Determining critical point on phase diagram

45. The Phase Diagrams of H 2 O and CO 2 Phase Diagrams

46. The Phase Diagrams of H 2 O and CO 2 Water: –The melting point curve slopes to the left because ice is less dense than water. –Triple point occurs at 0.0098  C and 4.58 mmHg. –Normal melting (freezing) point is 0  C. –Normal boiling point is 100  C. –Critical point is 374  C and 218 atm. Phase Diagrams

47. The Phase Diagrams of H 2 O and CO 2 Carbon Dioxide: –Triple point occurs at -56.4  C and 5.11 atm. –Normal sublimation point is -78.5  C. (At 1 atm CO 2 sublimes it does not melt.) –Critical point occurs at 31.1  C and 73 atm. Phase Diagrams

48. Prentice Hall © 2003Chapter 11 The forces holding solids and liquids together are called intermolecular forces. The covalent bond holding a molecule together is an intramolecular forces. The attraction between molecules is an intermolecular force. Intermolecular forces are much weaker than intramolecular forces When a substance melts or boils the intermolecular forces are broken (not the covalent bonds). Intermolecular Forces

49. Prentice Hall © 2003Chapter 11 Intermolecular Forces

50. Prentice Hall © 2003Chapter 11 Ion-Dipole Forces Interaction between an ion and a dipole (e.g. water). Strongest of all intermolecular forces. Intermolecular Forces

51. Prentice Hall © 2003Chapter 11 Dipole-Dipole Forces Exist between neutral polar molecules. Polar molecules need to be close together. Weaker than ion-dipole forces. There is a mix of attractive and repulsive dipole-dipole forces as the molecules tumble. If two molecules have about the same mass and size, then dipole-dipole forces increase with increasing polarity. Intermolecular Forces

52. Prentice Hall © 2003Chapter 11 Dipole-Dipole Forces Intermolecular Forces

53. Prentice Hall © 2003Chapter 11 Dipole-Dipole Forces Intermolecular Forces

54. Prentice Hall © 2003Chapter 11 London Dispersion Forces Weakest of all intermolecular forces. It is possible for two adjacent neutral molecules to affect each other. The nucleus of one molecule (or atom) attracts the electrons of the adjacent molecule (or atom). For an instant, the electron clouds become distorted. In that instant a dipole is formed (called an instantaneous dipole). Intermolecular Forces

55. Prentice Hall © 2003Chapter 11 London Dispersion Forces Polarizability is the ease with which an electron cloud can be deformed. The larger the molecule (the greater the number of electrons) the more polarizable. London dispersion forces increase as molecular weight increases. London dispersion forces exist between all molecules. London dispersion forces depend on the shape of the molecule. Intermolecular Forces

56. Prentice Hall © 2003Chapter 11 London Dispersion Forces The greater the surface area available for contact, the greater the dispersion forces. London dispersion forces between spherical molecules are lower than between sausage-like molecules. Intermolecular Forces

57. Prentice Hall © 2003Chapter 11 London Dispersion Forces Intermolecular Forces

58. Prentice Hall © 2003Chapter 11 London Dispersion Forces Intermolecular Forces

59. Prentice Hall © 2003Chapter 11 Hydrogen Bonding Special case of dipole-dipole forces. By experiments: boiling points of compounds with H-F, H-O, and H-N bonds are abnormally high. Intermolecular forces are abnormally strong. Intermolecular Forces

60. Prentice Hall © 2003Chapter 11 Hydrogen Bonding

61. Prentice Hall © 2003Chapter 11 Intermolecular Forces

62. Problems Chapter 11 2, 3, 15-19, 27, 49, 51, 53, 55, 67, 85, 87, 89, 90, 101, 102, 104