1. 11 Chapter 11 Intermolecular Forces, Liquids and Solids CHEMISTRY The Central Science 9th Edition David P. White

2. 21 11.1: A Molecular Comparison of Liquids and Solids

3. 31 The forces holding solids and liquids together are called intermolecular forces.

4. 41 The covalent bond holding a molecule together is an intramolecular forces The attraction between molecules is an intermolecular force Much weaker than intramolecular forces Melting or boiling: the intermolecular forces are broken (not the covalent bonds) 11.2: Intermolecular Forces

5. 51 The stronger the attractive forces, the higher the boiling point of the liquid and the melting point of a solid (low boiling point)

6. 61 Ion-Dipole Forces Interaction between an ion and a dipole (a polar molecule such as water) Strongest of all intermolecular forces (MIXTURES ONLY!)

7. 71 Dipole-Dipole Forces Between neutral polar molecules Oppositely charged ends of molecules attract Weaker than ion-dipole forces Dipole-dipole forces increase with increasing polarity Strength of attractive forces is inversely related to molecular volume

8. 81 London Dispersion Forces Weakest of all intermolecular forces Two adjacent neutral, nonpolar molecules The nucleus of one attracts the electrons of the other Electron clouds are distorted Instantaneous dipole Strength of forces is directly related to molecular weight London dispersion forces exist between all molecules Examples: 11.11 and 11.13

9. London dispersion forces depend on the shape of the molecule The greater the surface area available for contact, the greater the dispersion forces

10. 101 Hydrogen Bonding Special case of dipole-dipole forces H-bonding requires H bonded to an electronegative element (most important for compounds of F, O, and N)

11. Hydrogen Bonding Boiling point increases with increasing molecular weight. The exception is water (H bonding) Text, P. 413

12. Hydrogen Bonding

13. Solids are usually more closely packed than liquids (solids are more dense than liquids) Ice is ordered with an open structure to optimize H-bonding (ice is less dense than water)

14. Text, P. 417; example 11.9

15. 151 Sample Problems # 7, 9, 11, 13, 15, 17, 19

16. 161 Viscosity Viscosity is the resistance of a liquid to flow Molecules slide over each other The stronger the intermolecular forces, the higher the viscosity Viscosity increases with an increase in molecular weight 11.3: Some Properties of Liquids

17. 171 Surface Tension Surface molecules are only attracted inwards towards the bulk molecules Molecules within the liquid are all equally attracted to each other

18. 181 Surface tension is the amount of energy required to increase the surface area of a liquid –Cohesive forces bind molecules to each other (Hg) –Adhesive forces bind molecules to a surface (H 2 O) –If adhesive forces > cohesive forces, the meniscus is U-shaped (water in glass) –If cohesive forces > adhesive forces, the meniscus is curved downwards (Hg in barometer)

19. 11.4: Phase Changes Text, P. 420 (Endothermic) (Exothermic) (Endothermic) (Exothermic)

20. Generally heat of fusion (melting) is less than heat of vaporization (evaporation): it takes more energy to completely separate molecules, than to partially separate them Text, P. 420

21. 211 Heating Curves Plot of temperature change versus heat added is a heating curve During a phase change, adding heat causes no temperature change (equilibrium) –These points are used to calculate  H fus and  H vap

22. 221 Text, P. 421 Added heat increases the temperature of a consistent state of matter Energy used for changing molecular motion, no T change

23. 231 Critical Temperature and Pressure Gases are liquefied by increasing pressure at some temperature Critical temperature: the minimum temperature for liquefaction of a gas using pressure A high C.T. means strong intermolecular forces Critical pressure: pressure required for liquefaction

24. 241 Examples: # 31, 33, WDP # 48 Other WDP examples: # 44, 46, 50

25. 251 Explaining Vapor Pressure on the Molecular Level Some of the molecules on the surface of a liquid have enough energy to escape to the gas phase After some time the pressure of the gas will be constant at the vapor pressure (equilibrium) 11.5: Vapor Pressure

26. 261 Dynamic Equilibrium: the point when as many molecules escape the surface as strike the surface Vapor pressure is the pressure exerted when the liquid and vapor are in dynamic equilibrium Volatility, Vapor Pressure, and Temperature If equilibrium is never established then the liquid evaporates Volatile substances (high VP) evaporate rapidly The higher the T, the higher the average KE, the faster the liquid evaporates (hot water evaporates faster than cold water)

27. 271 Vapor pressure increases nonlinearly with increasing temperature Clausius-Clapeyron Equation

28. 281 When Temperature changes from T1 to T2, Pressure changes from P1 to P2 Use the Clausius-Clapeyron Equation to 1.Predict the vapor pressure at a specified temperature 2.Determine the T at which a liquid has a specified VP 3.Calculate enthalpy of vaporization from measurements of VP’s at different temperatures

29. 291 Sample problems: # 45, WDP # 35 Other WDP examples: # 36 & 37

30. 301 Vapor Pressure and Boiling Point Liquids boil when the external pressure equals the vapor pressure Normal BP: BP of a liquid at 1atmosphere Temperature of boiling point increases as pressure increases

31. 311 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 11.6: Phase Diagrams

32. 321 Text, P. 428 Vapor Pressure curve of the liquid (increase P, increase T) Stable at low P and high T Stable at low T and high P Triple Point: all 3 phases in equilibrium Beyond this point, liquid and gas phases are indistinguishable Melting point curve: Increased P favors solid phase; Higher T needed to melt the solid at higher P

33. 331 The Phase Diagrams of H 2 O and CO 2 Text, P. 429; Question 49 on P. 444 Line slopes to the left: ice is less dense than water (why?) MP decreases with increased P

34. 341 Sample Problem: BLBB #51

35. 351 Unit Cells Crystalline solid: well-ordered, definite arrangements of molecules, atoms or ions The smallest repeating unit in a crystal is a unit cell It has all the symmetry of the entire crystal Three-dimensional stacking of unit cells is the crystal lattice Close-packed structure 11.7: Structures of Solids

36. 361

37. Unit Cells

38. Primitive cubic: atoms at the corners of a simple cube each atom shared by 8 unit cells

39. 391 Unit Cells Body-centered cubic (bcc): atoms at the corners of a cube plus one in the center of the body of the cube corner atoms shared by 8 unit cells center atom completely enclosed in 1 unit cell

40. 401 Unit Cells Face-centered cubic (fcc): atoms at the corners of a cube plus one atom in the center of each face of the cube corner atoms shared by 8 unit cells face atoms shared by 2 unit cells

41. Unit Cells Text, P. 432 2 atoms per cell 4 atoms per cell 1 atom per cell

42. 421 The Crystal Structure of Sodium Chloride Two equivalent ways of defining unit cell: Cl- (larger) ions at the corners of the cell, or Na+ (smaller) ions at the corners of the cell

43. 431 11.8: Bonding in Solids Text, P. 435

44. 441 Covalent-Network Solids

45. Ionic Solids The structure adopted depends on the charges and sizes of the ions

46. 461 Metallic Solids Various arrangements are possible The bonding is too strong for London dispersion and there are not enough electrons for covalent bonds The metal nuclei float in a sea of electrons Metals conduct because the electrons are delocalized and are mobile Close-packed structure

47. 471 Amorphous solids (rubber, glass) have no orderly structure –IMFs vary in strength throughout the sample –No specific melting point

48. Sample Problems # 53, 69, 71, 73, 75

49. 491 End of Chapter 11 Intermolecular Forces, Liquids and Solids