ch12temp.pdf
- 1. 1 Chapter 12 Intermolecular Forces and the Physical Properties of Liquids and Solids Insert picture from First page of chapter Copyright McGraw-Hill 2009
- 2. Copyright McGraw-Hill 2009 2 12.1 Intermolecular Forces • Intermolecular forces are the attractive forces holding particles together in the condensed (liquid and solid) phases of matter • Result from coulombic attractions – Dependent on the magnitude of the charge – Dependent on distance between charges • Weaker than forces of ionic bonding • Involve partial charges
- 3. Copyright McGraw-Hill 2009 3 The Three Phases of Matter condensed phases
- 4. Copyright McGraw-Hill 2009 4 Types of Intermolecular Forces • van der Waals forces –between atoms and molecules of pure substances – Dipole-dipole interactions – attractive forces between polar molecules – Hydrogen bonding – attractive force in polar molecules containing a H atom bonded to a small, highly electronegative element (N, O and F) – (London) Dispersion forces – attractive forces arising from instantaneous dipoles and induced dipoles
- 5. Copyright McGraw-Hill 2009 5 Arrangement of Polar Molecules in a Liquid and a Solid.
- 6. Copyright McGraw-Hill 2009 6 • Intermolecular forces determine certain physical properties.
- 7. Copyright McGraw-Hill 2009 7 Hydrogen Bonds Between HF Molecules
- 8. Copyright McGraw-Hill 2009 8 Effect of Molar Mass and Hydrogen Bonding on Boiling Points
- 9. Copyright McGraw-Hill 2009 9 Instantaneous and Induced Dipoles Magnitude depends on the ability to be polarized which is greater for larger molecules.
- 10. Copyright McGraw-Hill 2009 10 Polarization and Molar Mass
- 11. Copyright McGraw-Hill 2009 11 What kind(s) of intermolecular forces exist in CH2ClCH2COOH(l) C C H Cl H C H H O O H dispersion forces dipole-dipole interactions hydrogen bonding
- 12. Copyright McGraw-Hill 2009 12 –Occur in mixtures of ionic and polar species –Coulombic attraction between ions and polar molecules –Dependent upon • Size and charge of ion • Dipole moment of the molecule • Size of the molecule –Can also be repulsive •Ion-dipole Interactions
- 13. Copyright McGraw-Hill 2009 13 12.2 Properties of Liquids • Surface Tension – a quantitative measure of the elastic force at the surface of a liquid • Manifestations – Formation of a mensicus – Capillary action which results from a combination of • Cohesion (attractions between like molecules, cohesive forces) • Adhesion (attractions between unlike molecules, adhesive forces)
- 14. Copyright McGraw-Hill 2009 14 Effect of Surface Tension
- 15. Copyright McGraw-Hill 2009 15 Intermolecular Forces: Surface versus Interior of a Liquid
- 16. Copyright McGraw-Hill 2009 16 Cohesion and Adhesion Water Mercury Adhesion > Cohesion Cohesion > Adhesion
- 17. Copyright McGraw-Hill 2009 17 • Viscosity – a measure of a fluid’s resistance to flow –Units: N.s/m2 –The higher the viscosity the greater the resistance to flow –Varies inversely with temperature –Stronger intermolecular forces produce higher viscosities
- 18. Copyright McGraw-Hill 2009 18 Glycerol – high viscocity due to • Three hydrogen bonding sites • Molecular shape
- 19. Copyright McGraw-Hill 2009 19 • Vapor Pressure of a Liquid –Depends on the magnitude of intermolecular forces –Temperature dependent T1<T2
- 20. Copyright McGraw-Hill 2009 20 • Equilibrium Vapor Pressure – a dynamic state − Evaporation (vaporization) – liquid molecules escape into the gas phase − Condensation – gas molecules return to the liquid phase
- 21. Copyright McGraw-Hill 2009 21 Comparison of Rates of Evaporation and Condensation at Constant Temperature
- 22. Copyright McGraw-Hill 2009 22 Effect of Temperature and Intermolecular Forces on Vapor Pressure
- 23. Copyright McGraw-Hill 2009 23 • Clausius-Clapeyron Equation – linear relation between temperature and vapor pressure − At two temperatures, T1 and T2: where R = 8.314 J/K . mol
- 24. Copyright McGraw-Hill 2009 24 An unknown compound exhibits a vapor pressure of 255 mmHg at 25.5oC and 434 mmHg at 48.8oC. What is DHvap of this substance?
- 25. Copyright McGraw-Hill 2009 25 K 298.65 273.15 C 25.5 1 o T K 321.95 273.15 C 48.8 2 o T D K 298.65 1 K 321.95 1 mol 8.314J/K mmHg 434 mmHg 255 vap H ln 3 3 3.3484x10 3.1061x10 mol 8.314J/K 0.53178 D vap H D 1 2 vap 2 1 1 1 T T R H P P ln
- 26. Copyright McGraw-Hill 2009 26 vap H D 1 4 K 2.423x10 mol) .314J/K 0.53178)(8 ( 1 4 K 2.423x10 mol 8.314J/K 0.53178 D vap H vap 4 J/mol 10 x 1.82 H D vap kJ/mol 18.2 H D
- 27. Copyright McGraw-Hill 2009 27 12.3 Crystal Structure • Crystalline solid – possesses rigid and long-range order • Lattice structure – arrangement of particles in a crystalline solid – Depends on nature of particles – Depends on size of particles • Stability depends on type of force between particles (ionic or covalent bonds and/or intermolecular forces)
- 28. Copyright McGraw-Hill 2009 28 unit cell – basic repeating structural unit of a crystalline solid lattice point – each atom ion or molecule in a unit cell single unit cell 3-D array of unit cells
- 29. Copyright McGraw-Hill 2009 29 Seven Types of Unit Cells
- 30. Copyright McGraw-Hill 2009 30 Coordination number – number of atoms (particles) surrounding an atom in a crystal lattice • Indicates how tightly atoms pack • Larger coordination numbers indicate tighter packing
- 31. Copyright McGraw-Hill 2009 31 • Types of cubic unit cells –simple or primitive (scc) –body-centered (bcc) –face-centered (fcc)
- 32. Copyright McGraw-Hill 2009 32 Alternate Perspective of bcc Arrangement
- 33. Copyright McGraw-Hill 2009 33 Sharing of Atoms by Adjacent Unit Cells corner atom edge atom face atom
- 34. Copyright McGraw-Hill 2009 34 • Number of atoms per unit cell –scc: 1 atom –bcc: 2 atoms –fcc: 4 atoms •Allocation of atoms among unit cells − corner atoms – 1/8 atom within unit cell − face atoms – 1/2 atom within unit cell − body atoms – 1 atom within unit cell
- 35. Copyright McGraw-Hill 2009 35 • Closest Packing – most efficient way to arrange atoms in a crystal –hexagonal closest packed (ABA) –cubic closest packed (ABC)
- 36. Copyright McGraw-Hill 2009 36 Closest Packing and Cubic Unit Cells fcc cubic hexagonal
- 37. Copyright McGraw-Hill 2009 37 Geometric Relationships
- 38. Copyright McGraw-Hill 2009 38 When silver crystallizes, it forms face- centered cubic cells. The unit cell edge length is 4.087A . Calculate the density of silver. o
- 39. Copyright McGraw-Hill 2009 39 g 10 7.167 amu 10 6.022 g 1 atom amu 107.9 cell unit atoms 4 22 23 m 3 23 3 8 3 cm 10 6.827 cm 10 4.087 a V cm 10 4.087 m 1 cm 00 1 A 10 1 m 1 A 4.087 8 10 a 3 3 23 22 g/cm 10.5 cell /unit cm 10 6.827 cell g/unit 10 7.167 V m d Mass of unit cell Volume of unit cell Density of unit cell
- 40. Copyright McGraw-Hill 2009 40 • X-ray diffraction utilizes the scattering of X-rays and the resulting scattering patterns to deduce arrangement of particles
- 41. Copyright McGraw-Hill 2009 41 − Bragg equation
- 42. Copyright McGraw-Hill 2009 42 12.4 Types of Crystals • Ionic Crystals – Composed of anions and cations – Held together by coulombic forces – Anions generally are bigger than cations – Size and relative number of each ion determines the crystal structure
- 43. Copyright McGraw-Hill 2009 43 Unit cell of NaCl as Defined by Cl− or Na+
- 44. Copyright McGraw-Hill 2009 44 Examples of Ionic Crystal Lattices CsCl ZnS CaF2
- 45. Copyright McGraw-Hill 2009 45 How many of each ion are contained within a unit cell of CaF2?
- 46. Copyright McGraw-Hill 2009 46 Ca2+ 8 corner ions x 1/8= 1 ion 6 face ions x ½ = 3 ions 4 ions of Ca2+ F− 8 body ions x 1 = 8 ions of F− Ca2+ F−
- 47. Copyright McGraw-Hill 2009 47 • Covalent crystals – Held together by covalent bonds diamond graphite
- 48. Copyright McGraw-Hill 2009 48 • Molecular crystals – Lattice points occupied by molecules – Held together by intermolecular forces water
- 49. Copyright McGraw-Hill 2009 49 • Metallic crystals – Lattice points occupied by atoms – Generally bcc, fcc, hexagonal closest packed – Very dense – Bonding arises from delocalized electrons
- 50. Copyright McGraw-Hill 2009 50
- 51. Copyright McGraw-Hill 2009 51 12.5 Amorphous Solids • Lack regular arrangement of atoms • Glass is a familiar and important amorphous solid – Transparent fusion of inorganic materials – Chief component SiO2 – Behaves more as a liquid than a solid
- 52. Copyright McGraw-Hill 2009 52 Comparison of crystalline quartz and amorphous quartz glass crystalline amorphous
- 53. Copyright McGraw-Hill 2009 53
- 54. Copyright McGraw-Hill 2009 54 12.6 Phase Changes • Phase – homogenous part of a system that is separated from the rest of the system by a well-defined boundary • Phase change – transition from one phase to another –Caused by the removal or addition of energy –Energy involved is usually in the form of heat
- 55. Copyright McGraw-Hill 2009 55 The Six Possible Phase Changes
- 56. Copyright McGraw-Hill 2009 56 • Liquid-Vapor Phase Transition – Boiling point – the temperature at which the vapor pressure of liquid equals atmospheric pressure – Molar heat of vaporization (DHvap) – the amount of heat required to vaporize one mole of a substance at its boiling point usually in kJ/mol – Dependent on the strength of intermolecular forces – Condensation – opposite of vaporization
- 57. Copyright McGraw-Hill 2009 57
- 58. Copyright McGraw-Hill 2009 58 • Critical temperature (Tc) – the temperature above which a gas cannot be liquified by application of pressure • Critical pressure (Pc) – the pressure that must be applied to liquefy a gas at Tc. • Supercritical fluid – the fluid that exists above Tc and Pc.
- 59. Copyright McGraw-Hill 2009 59
- 60. Copyright McGraw-Hill 2009 60 • Solid-Liquid Phase Transition – Freezing – transformation of liquid to solid – Melting (fusion) – opposite of freezing – Melting point of solid (or freezing point of liquid) – temperature at which the solid and liquid phases coexist in equilibrium • Dynamic equilibrium in which the forward and reverse processes are occurring at the same rate – Molar heat of fusion (DHfus) – energy to melt one mole of a solid usually in kJ/mol
- 61. Copyright McGraw-Hill 2009 61 Typical Heating Curve
- 62. Copyright McGraw-Hill 2009 62 • Solid-Vapor Phase Transition –Sublimation – process by which molecules go directly from the solid phase to the vapor phase –Deposition – reverse of sublimation –Molar heat of sublimation (DHsub) – energy required to sublime one mole of solid usually in kJ/mol iodine
- 63. Copyright McGraw-Hill 2009 63 Heating Curve for Water
- 64. Copyright McGraw-Hill 2009 64 Calculate the amount of energy (in kJ) required to convert 125 g of ice at 10.0oC to liquid water at the normal boiling point. Assume that the specific heat of ice is 2.050 J/goC.
- 65. Copyright McGraw-Hill 2009 65 Energy to warm ice from 10oC to 0oC C 10.0 C) 10.0 ( C 0.0 o o o DT kJ 10 2.563 C 10.0 C g J 2.050 g 125 3 o o D T ms q Energy to melt ice at 0oC kJ 10 4.169 mol kJ 6.01 mol 937 . 6 1 vap D H n q kJ 2.563 J 10 1 kJ J 10 2.563 3 3 mol 6.937 g 18.02 mol g 125
- 66. Copyright McGraw-Hill 2009 66 Energy to warm water from 0.0oC to 100.0oC C 100.0 C 0.0 C 100.0 o o o DT J 10 5.230 C 100.0 C g J 4.184 g 125 4 o o D T ms q Total energy required kJ 10 5.230 J 10 1 kJ 1 J 10 5.230 1 3 4 96.6kJ ) kJ 10 5.230 ( ) kJ 10 4.169 ( kJ 2.563 1 1
- 67. Copyright McGraw-Hill 2009 67 12.7 Phase Diagrams • Phase diagram – summarizes the conditions (temperature and pressure) at which a substance exists as a solid, liquid or gas – Divided into three regions (solid, liquid, gas) – Phase boundary line – line separating any two regions – Triple point – the point at which all three phase boundary lines meet
- 68. Copyright McGraw-Hill 2009 68 Phase Diagram of CO2
- 69. Copyright McGraw-Hill 2009 69 Heating CO2 Starting at 100oc and 1 atm
- 70. Copyright McGraw-Hill 2009 70 Phase Diagram of H2O
- 71. Copyright McGraw-Hill 2009 71 What is a) the normal* melting point, b) the normal* boiling point and c) the physical state of the substance at 2.0 atm and 110o C? *normal – measured at 1.00 atm
- 72. Copyright McGraw-Hill 2009 72 normal melting point normal boiling point ~135oC ~200oC
- 73. Copyright McGraw-Hill 2009 73 solid physical state
- 74. Copyright McGraw-Hill 2009 74 Key Points • Intermolecular forces – Dipole-dipole interactions – Hydrogen bonding – (London) dispersion forces • Properties of liquids – Surface tension – Viscosity – Vapor pressure • Clausius-Clapeyron equation
- 75. Copyright McGraw-Hill 2009 75 • Crystal structure – Unit cells • Lattice point – Packing spheres • Coordination number • Cubic unit cells – Closest Packing • Types of crystals – Ionic – Covalent – Molecular – Metallic
- 76. Copyright McGraw-Hill 2009 76 • Amorphous solids • Phase changes – Liquid-vapor transitions • Boiling point • Heat of vaporization • Critical temperature and pressure – Solid-liquid transitions • Melting point • Heat of fusion • Phase diagrams