1. Intermolecular Forces, Liquids, and Solids Chapter 11/12

2. Intermolecular Forces vs. Intramolecular Forces

4. Intermolecular Forces The stronger the attractive forces in a liquid, the higher the temperature at which the liquid boils

5. Intermolecular Forces The stronger the attractive forces in a solid, the higher the temperature at which the solid melts

6. Three Types of Intermolecular Forces Dipole-Dipole Force, London Dispersion forces, and hydrogen bonding forces A fourth force is an Ion- Dipole Force

7. Forces All four forces are electrostatic and tend to be less than 15% as strong as an ionic or covalent bond

8. Ion-Dipole Forces exist between the ion and the partial charge on the end of a polar molecule important in solutions

10. Dipole-Dipole Forces Neutral polar molecules attract to each other when the positive end of one molecule is near the negative end of another

11. Dipole-Dipole Forces effective when polar molecules are close to one another weaker than ion-dipole forces

12. Dipole-Dipole Forces For molecules of equal mass and size, the strengths of intermolecular attractions increase with increasing polarity

14. London Dispersion Forces the motion of electrons in an atom or molecule can create an instantaneous dipole moment

15. London Dispersion Forces In one second, when the electrons are on the same side of an atom, it can influence its neighbor causing an attraction between the two

16. London Dispersion Forces The ease at which the charge distribution in a molecule can be distorted is called polarizability

17. London Dispersion Forces The more polarizability in a molecule, the stronger the London dispersion forces

18. London Dispersion Forces increase in strength with increasing molecular weight and depend on shape exist between all molecules, polar and nonpolar

20. Sample Exercise 11.1 Comparing Intermolecular Forces The dipole moments of acetonitrile, CH 3 CN, and methyl iodide, CH 3 I, are 3.9 D and 1.62 D, respectively. (a)Which of these substances has greater dipole– dipole attractions among its molecules? (b)Which of these substances has greater London dispersion attractions? (c)The boiling points of CH 3 CN and CH 3 I are 354.8 K and 315.6 K, respectively. Which substance has the greater overall attractive forces?

21. Sample Exercise 11.1 Comparing Intermolecular Forces Of Br 2, Ne, HCl, HBr, and N 2, which is likely to have (a) the largest intermolecular dispersion forces, (b) the largest dipole–dipole attractive forces?

22. Hydrogen Bonding intermolecular attraction between the H atom in a polar bond(H-F, H-O, or H-N) and an unshared electron pair on a nearby electronegative ion or atom(usually N,F, or O)

23. Hydrogen Bonding Considered a unique dipole- dipole interaction due to the electronegativity of N,F, and O Stronger than dipole-dipole or dispersion forces

24. Hydrogen Bonding

25. Identify the attractive forces between the following:

26. Sample Exercise 11.2 Identifying Substances that Can Form Hydrogen Bonds In which of the following substances is hydrogen bonding likely to play an important role in determining physical properties: CH 4, H 2 NNH 2, CH 3 F, H 2 S,CH 2 Cl 2, H 2 O 2, CH 3 COCH 3, PH 3 ?

27. Comparing Intermolecular Forces

28. Sample Exercise 11.3 Predicting the Types and Relative Strengths of Intermolecular Attractions List the substances H 2, CO, HF, and Ne in order of increasing boiling points. Identify IMF’s present in the following substances and order their bp: CH 3 CH 3, CH 3 OH, and CH 3 CH 2 OH.

29. Properties of Liquids Viscosity is the resistance of a liquid to flow Related to molecule movement

30. Properties of Liquids Greater viscosity results from greater attractive forces which slow the flow Viscosity decreases with increasing T

31. Viscosity

32. Properties of Liquids Surface tension is the energy required to increase the surface area of a liquid by a unit amount. Results from imbalance of IMF’s on surface experiencing net inward force

33. Surface tension

34. Phase Changes

35. accompanied by a change in energy of the system

36. Heating Curves Graph of the temperature of the system versus the amount of heat added.

38. Phase Diagrams graphical way to summarize the conditions under which equilibrium exists between different states of matter

39. General Phase Diagram Includes three important curves which represent temperatures and pressures where the various phases can co-exist

41. General Phase Diagram The line from A to B is the vapor pressure curve. It represents the equilibrium between the liquid and gas phase. The normal bp is found at 1 atm on this curve

42. Critical Point Point B is where the critical P and T are found Beyond the critical point the liquid and gas phase are indistinguishable

43. General Phase Diagram Line AC represents the equilibrium between solid and gas

44. General Phase Diagram The line AD represents the equilibrium between the solid and liquid phase The mp and fp are identical, and differ only in direction The normal mp is at 1 atm

45. General Phase Diagram Point A is the Triple Point, where the three phases intersect and are at equilibrium

46. Phase Diagram of Water the mp of water decreases with increasing pressure because the liquid form is more compact than the solid form

47. Phase Diagram of CO 2 There is no normal mp in CO 2. At 1 atm pressure, the substance will sublime

49. Sample Exercise 11.6 Interpreting a Phase diagram Describe any changes in the phases present when H2O is (a)kept at 0 °C while the pressure is increased from that at point 1 to that at point 5 (b) kept at 1.00 atm while the temperature is increased from that at point 6 to that at point 9

50. Structures of Solids Crystalline Solids –Atoms, ions, or molecules are ordered in a well-defined three-dimensional arrangements; they have flat surfaces or faces and regular shapes –Quartz and diamond are crystalline solids

51. Structures of Solids Amorphous Solids –Solids in which particles have no orderly structure; lack faces and shape –Rubber and glass are amorphous solids

52. Structures of Solids

53. Bonding in Solids - Molecular Solids Atoms or molecules (nonmetals) held together by IMF’s Soft, low mp, poor conductors

54. Bonding in Solids – Covalent-Network Solids Atoms held together in large networks or chains by covalent bonds Hard, high mp, variable conductors

55. Bonding in Solids - Ionic Solids Ions held together by ionic bonds (electrostatic forces) Hard, brittle, high mp, poor conductors

56. Bonding in Solids - Metallic Solids Metal atoms Visualize the metals as an array of positive ions immersed in a sea of delocalized valence electrons soft to very hard, low to high mp, excellent conductors, malleable and ductile

57. Alloys Mixture of elements exhibiting metallic properties

58. Summary of Bonding in Solids