Chapter 10: Intermolecular Forces, Liquids, and Solids

1. Chapter 10: Intermolecular Forces, Liquids, and Solids

2. Ch. 10 A Molecular Comparison of Liquids and Solids • Gases characterized by rapidly moving, widely-spaced particles • Solids characterized by a regular array of closely-spaced, fixed particles • Liquids are somewhere in-between, but with some special properties of their own • Why do liquids and solids exist? Why is all matter not in the gaseous state? 01m07an1

3. States of Matter

4. Changes of State • Converting a gas into a liquid or solid requires the molecules to get closer to each other: • cool or compress. • Converting a solid into a liquid or gas requires the molecules to move further apart: • heat or reduce pressure. • The forces holding solids and liquids together are called intermolecular forces.

5. Whiteboards • Helium • Melting Point: 0.95 K (@ 2.5 MPa, about 24.6 times greater than standard pressure 101.5 kPa) • Boiling Point: 4.22 K

6. London Dispersion Forces • London dispersion forces are present between all molecules; • temporary uneven charges give a temporarily polar molecule • Temporarily polar molecule influences / polarizas molecules nearby London4

7. a Attraction occurs through the outer unoccupied orbitals of atoms/molecules. Referred to as a Van der Waals force or london dispersion force (NOT a bond)

8. London Forces • The magnitude of London forces explains why Cl2 is a gas, Br2 is a liquid, and I2 is a solid. • Explain the order of the boiling points of the halogens and noble gases:

9. London Dispersion Forces • Only force between non-polar molecules or between non-bonded atoms (group 8A) • London forces become stronger, the larger the atom or molecule • Strength increases with polarizability • larger electron clouds are easier to deform, “squishiness” of the electron cloud • Found in mixtures or pure substances.

10. Boiling Points

11. Boiling Points • Boiling point increases with the size of molecules because of increases in London forces with larger electron clouds.Figure 10.3

12. London Dispersion Forces • Which member of each pair has the stronger London forces? Ne, Kr F2, Cl2 CH4, SiCl4 N2, O2 • N2 b.p. = 77.4 K • O2 b.p. = 90.2 K

13. London Forces in Molecular Solids • Fig 10.34 in Ch 10: • Sulfur, S8 • Phosphorus, P4 • Different allotropes will have different properties based on the structural differences • Black, red, and white phosphorus

14. London Forces in Organics • Examples: • methane • propane • pentane • Rank the compounds in order of decreasing boiling point. Can you explain the trend?

15. London Forces in Organics • Whiteboard structural formulas for: Isomer B.P. • methane -162 oC • propane -42 oC • hexane 68 oC • Explain the trend in boiling points!

16. London Forces in Organics • Whiteboard structural formulas for: Isomer B.P. • Pentane 36.1 oC • 2-methylbutane 27.7 oC • 2,2-dimethylpropane 9.5 oC • Explain the trend in boiling point

17. Whiteboards • Hydrogen Chloride • Melting Point: 159 K • Boiling Point: 188 K • Compare HCl to Argon (a noble gas with similar molar mass): • Argon, Ar • Melting Point: 83.8 K • Boiling Point: 87.3 K • What can account for these differences?

18. Dipole-Dipole Forces • Molecules that are permanent dipoles attract each other’s polar ends. 13m05an2 dipole 1 Stirring bar demo

19. Dipole-dipole attractions (431 kJ/mol) (16 kJ/mol) Still NOT bonds!!!

20. Dipole-Dipole Forces • Permanent dipoles attract one another • Pure substance or mixture • Strength of these forces depend on the dipole moment of the molecules 

21. Dipole-Dipole Forces • Explains why polar liquids are more soluble in polar liquids than in non-polar liquids • It takes 2000 mL of H2O to dissolve 1 mL of CCl4 • It takes 50 mL of H2O to dissolve 1 mL of CH2Cl2 • Which member of each pair has the stronger intermolecular forces? SiCl4, SiHCl3 CO2, SO2

22. Ion-Dipole Attractions • Ions have full charges that are attracted to the partial charge on polar molecules (dipoles) • Helps explain solubility of ionic salts in polar solvents (like water)

23. Ion-Dipole Attractions • Metal ions are hydrated in solution; this hydrate often is found in solid metal salts as well. Some hydrates actually form coordinate-covalent bonds and are very stable. • Fe(H2O)63+ is an example • Cr(H2O)63+ is so stable that it takes many hours to exchange a bonded water molecule for one in the solvent water

24. Hydrated Iron(III) Ion • The water molecules are attracted so strongly in this case that they donate electron pairs to the iron ion and form a coordinate covalent bond.

25. Strange Behavior… • On a whiteboard, draw Lewis structures for: • Water, H2O • Ethanol, C2H5OH • Acetone, C3H6O • Identify the intermolecular forces present in a sample of each substance. Which sample should evaporate fastest? Why? • Observe the demonstration on your whiteboard!

26. Strange Behavior… • Compare observations to heats of vaporization of these substances. molar mass ΔHvap • Water: 18.0 g/mol 40.7 kJ/mol • Ethanol: 58.0 g/mol 38.6 kJ/mol • Acetone: 70.0 g/mol 31.3 kJ/mol Whiteboards: • How can our model of interacting particles account for these differences?

27. Strange behavior… • Trend in boiling points of polar moleculesalso reveals an especially strong dipole-dipole force:

28. Hydrogen Bonding in Liquid Water • H points at the electron pair on the atom in the other molecule • In liquid water, each water molecule is surrounded by an average of 4 other water molecules; structure is not rigid. • Longer than covalent bond.

29. Hydrogen Bonding • Average of 4 hydrogen bonds in liquid waterFigure 10.2 • Fluoride ion is hydrogen-bonded to water in solution

30. Structure of Ice • The water molecules in ice are fixed into a tetrahedral arrangement. Open structure makes ice less dense than water.

31. Structure of Ice • The open structure ofice leaveschannels ofempty spacethrough thecrystals.

32. Hydrogen Bonding • Molecules hydrogen-bond to themselves or to other molecules.Figure 10.2

33. Hydrogen Bonding • H-bonding is observed for HF, H2O, NH3, but not CH4 • Conditions for occurrence: • H attached to a small, highly electronegative element in one molecule • Small, highly electronegative element with one or more unshared electron pairs in the other molecule • Observed for the elements: F, O, N(rarely S and Cl)

34. Intermolecular Forces • What types of intermolecular forces are observed for each of the following molecules?(A molecule may have more than one.) H2O HF HBr NH3 PF3 CH3OH F2 CO CO2 N2

35. Trends in Intermolecular Forces • Which member of each pair has the larger intermolecular forces (and therefore boiling point, heat of vaporization)? CH3OH, CH3SH F2, Kr F2, CO CO, HF CO2, NH3 N2, NH3

36. Summarize! Type and Strengths of Intermolecular Forces • Intermolecular forces generally increase in strength as London < Dipole-Dipole < H-bonding < Ion-Dipole attractions (mixtures) < Ionic Bonding / network-covalent / metallic • The forces are cumulative. All molecules have London forces. Polar molecules have both London and dipole-dipole forces. ...

37. Identify PredominantType of Intermolecular Forces

38. Warm-Up Quiz • Half sheet of paper, you may use notes: • Which member of each pair has the larger intermolecular forces (boiling point, heat of vaporization)? • CH4, CH3CH3 • NH3, PH3 • CO2, SO2 • CO, CO2 • I2, Cl2

39. How does a gecko climb?

40. Importance of Intermolecular Forces • Intermolecular forces (forces between molecules) are weaker than bonds, but have profound effects on the properties of liquids • Polar liquids have a higher boiling point and higher heat of vaporization than non-polar liquids. • Polar liquids dissolve ionic solids and polar liquids.

41. 10.2 The Liquid State • Intermolecular forces and distance between particles are intermediate. • Highly disordered structure like gases, but closely packed particles like solids. • Properties intermediate between those of gases and solids

42. Viscosity • Viscosity – measure of a liquid’s resistance to flow • Values depend on molecular size and intermolecular forces • Water is about average (high intermolecular forces, but small size) • How could viscosity be measured in the lab? 13m11vd2

43. Measuring Viscosity • Techniques for measuring viscosity in the lab: • Drain method: time it takes for a fluid to flow through an opening (i.e. out of a buret or pipet) • “Falling ball” method: time it takes for a ball or other object to fall through a liquid. • More information available at: • Techniques for Measuring Viscosity in the Field

44. Surface Tension How does a water strider stay on the top of the water? Why does the needle float?

45. Surface Tension • Surface tension results from the amount of energy required to increase the surface area by a given amount. • No counterpart in gases or solids • Caused by intermolecular forces, which are larger in the bulk of the liquid than at the surface. • Molecules must break IM forces in order to move to the surface and increase the surface area.

46. Surface Tension • Why are water droplets spherical?

47. Measurement of Surface Tension • Surface tension, at least for comparative purposes, can be studied by examining droplets of water. • “Penny drops” – examine the number of drops of liquid that will stay on the surface of a penny. • Important to repeat trials to check precision. • Can compare multiple liquids. • How does number of drops on penny relate to surface tension?

48. Capillary Action • Why do liquids rise in a capillary? 13m11an1

49. Capillary Action • Which forces are responsible for these observations? solids

50. Capillary Action • Competing forces: • Between liquid molecules (cohesive) • Between liquid molecules and glass (adhesive) • H2O-Glass > H2O-H2O • Hg-Glass < Hg-Hg • Height determined by net liquid-glass forces opposing gravityFigure 10.8