1 Chapter 11 Fritz London Studied intermolecular induced-dipole interactions. Johannes D. van der Waals * Studied intermolecular forces in VPT relationships in liquids and gases. Liquids, Solids, and Materials

2 Kinetic Energy Attractive Intermolecular Forces High temperatures Low temperatures Gases, Liquids and Solids

3 Ionic Forces Ion-Ion e.g. NaCl(s) Ion-Dipole e.g. NaCl(aq) Ions form strong intermolecular forces with the polar molecule water. Ion-ion forces are very strong and produce high boiling points and melting points. NaCl dissolved in water From now on we will concentrate on covalent molecules

4 Intermolecular Interactions Non-Polar Molecules: Polar Molecules: Molecules with F-H, O-H, N-H Bonds: 1. Dispersion (Disp) Forces 2. Dipole-Dipole (DD) Forces 1. Dispersion (Disp) Forces 2. Dipole-Dipole (DD) Forces 3. Hydrogen Bonding (HB) Forces

5 London Dispersion Forces (aka van der Waals Forces) The electrons on one atom are attracted to the nucleus on a neighboring atom. This creates an “instantaneous” (i.e. temporary) dipole on the first atom. The instantaneous dipole on the first atom then induces an instantaneous dipole on the second atom. The two induced dipoles attract each other. London Dispersion Forces are proportional to a molecule’s polarizability, which is the ease with which the electron cloud can be deformed. The polarizability, is approximately proportional to the number of electrons in the molecule.

6 Element MW # e - T bp [amu] [K] He Ne Ar Kr Xe The noble gases are spherical, non-polar atoms. As the molecule (atom) increases in size, the boiling point increases (vapor pressure decreases). Noble Gases

7 London Dispersion Forces London Dispersion Force  # e - London Dispersion Force  MW T bp is a good measure of the strength of intermolecular forces. A higher T bp indicates stronger intermolecular forces. Compd. MW # e - T bp [amu] [ o C ] F Cl Br I Compd. MW # e - T bp [amu] [ o C ] CH C 2 H C 3 H n-C 4 H Dispersion Forces Increasing

8 Dispersion forces also depend upon the molecular shape. Pentane: C 5 H 12 - MW = 72 amu n-pentane T bp = 36 o C neopentane T bp = 9 o C Large Contact Area (Surface Area) Small Contact Area (Surface Area)

9 Dipole-Dipole Forces In liquids of polar molecules, oppositely charged ends of the molecules tend to attract each other, causing partial alignment. For molecules of roughly equal MW’s (i.e. with similar Dispersion Forces), the molecule with the higher dipole moment will have a higher boiling point due to greater Dipole-Dipole forces.

10 Propane CH 3 CH 2 CH 3 Acetonitrile CH 3 CN MW = 44 amu T bp = -42 o C MW = 41 amu T bp = +82 o C   0 No Dipole Moment  = 3.9 Debye (D) Large Dipole Moment Why is the boiling point of acetonitrile so much higher than the boiling point of propane??? Dipole-Dipole Forces Consider this….

11 Hydrogen Bonding F-H Bonds: H-F O-H Bonds: e.g. H 2 O, CH 3 OH N-H Bonds: e.g. NH 3, CH 3 NH 2 H bonding: between H and a very electronegative atom

12 Compd. MW # e -  [amu] [D] F-F H-Cl T bp [ o C ] H-F Why so high?? Hydrogen Bond ++ -- ++ --

13 Hydrogen Bonding Compd. MW # e - [amu] H 2 S H 2 O T bp [ o C ] PH NH

14 CH 3 CH 2 OH HOCH 2 CH 2 OH CH 3 OCH 3 CH 3 CH 2 CH 3 MW=74 MW=62 MW=46 MW=44 Disp Disp Disp Disp DD DD DD HB Lowest BP Highest BP Ethylene glycol - a viscous liquid; keeps radiator fluid from boiling over in your car (198°) Propane - a gas used for charcoal Grills (- 42°) Ether - a very volatile Liquid (35°) Ethyl alcohol - a liquid with slightly lower bp than water (78°)

15 Vapor Pressure Vap. Press. (H 2 O) liquid water H2OH2O H2OH2O H2OH2O H2OH2O H2OH2O H2OH2O H2OH2O H2OH2O H2OH2O H2OH2OH2OH2O H2OH2O

16 Vapor Pressure Rises with Temperature Temperature Vapor Pressure At higher temperatures, more molecules have sufficient energy to escape from the liquid.

17 The Boiling Point The boiling point of a liquid is the temperature at which its vapor pressure is equal to the external pressure. The “normal” boiling point is the temperature at which the vapor pressure of the liquid is equal to 760 torr (1 atm.)

18 The Critical Temperature H2OH2O T Vap. Press. [ o C] [atm]  T c = Critical Temperature Highest temperature at which substance can be liquified. P c = Critical Pressure Pressure required to liquify substance at T c T c = 374 o C P c = 218 atm

19 Substance T c P c Helium -268 o C 2.3 atm Oxygen Ethane Propane Freon (CCl 2 F 2 ) Water The Critical Temperature

20 Phase Diagrams 1. Why can’t you ice skate in Red Lake, Minnesota in February? 2. Why does carbon dioxide sublime, whereas water first melts and then vaporizes? to St. Paul

21 Phase Diagram of H 2 O Temperature ( o C) Pressure (atm) SolidLiquidVapor

22 A D B C Temperature ( o C) Pressure (atm) SolidLiquidVapor AB: Liquid-Vapor Equilibrium AC: Solid-Liquid Equilibrium AD: Solid-Vapor Equilibrium A: Triple Point Solid-Liquid-Vapor B: Critical Point torr Phase Diagram of H 2 O

23 Pressure dependence of T bp Temperature ( o C) Pressure (atm) LiquidVapor Liquid  Vapor d(liq) >> d(vap) V(liq) << V(vap) Increased pressure shifts equilibrium in direction of lower volume.

24 Pressure dependence of T mp in H 2 O Temperature ( o C) Pressure (atm) Solid (ice)  Liquid d(ice) < d(liq) V(ice) > V(liq) Increased pressure shifts equilibrium in direction of lower volume. SolidLiquid Application to ice skating

25 P1P1 Temperature ( o C) Pressure (atm) SolidLiquidVapor P2P2 Melting point decreases with Pressure Melting + Vaporization Sublimation 4.6 torr 0.01 Phase Diagram of H 2 O

26 Phase Diagram of “Normal” Substances Temperature ( o C) Pressure (atm) P1P1 SolidLiquidVapor P2P2 Melting point increases with Pressure Melting + Vaporization Sublimation CO

27 Pressure dependence of T mp in “normal” substances Temperature ( o C) Pressure (atm) 1100 Solid  Liquid d(sol) > d(liq) V(sol) < V(liq) Increased pressure shifts equilibrium in direction of lower volume. SolidLiquid T mp o T mp

28 Phase Transitions Solid Gas Liquid Melting Freezing Vaporization Enthalpy SublimationDeposition Condensation

29 Energy Changes of Phase Transitions Solid Liquid Melting Freezing Enthalpy Melting (Fusion)  H fus = H liq - H sol = 6.01 kJ/mol (for H 2 O) ENDOTHERMIC Freezing (Crystallization)  H crys = H sol - H liq = kJ/mol (for H 2 O) EXOTHERMIC

30 Gas Liquid Vaporization Enthalpy Condensation Vaporization  H vap = H gas - H liq = 40.7 kJ/mol (for H 2 O) ENDOTHERMIC Condensation  H cond = H liq - H gas = kJ/mol (for H 2 O) EXOTHERMIC Energy Changes of Phase Transitions

31 Heating Curves Heat Added (Joules) Temperature ( o C) T mp T bp 1. Heating solid 1 2. Melting solid to liquid 2 3. Heating liquid 3 4. Vaporizing liquid to gas 4 5. Heating gas 5

32 Quantitative considerations of heating curves  H fus = 6.0 kJ/mol  H vap = 40.7 kJ/mol T mp = 0 o C T bp = 100 o C Heat capacities: C s (sol)= 2.09 J/g- o C C s (liq)= 4.18 J/g- o C C s (gas)= 1.84 J/g- o C For water: To heat 18 g of H 2 O from -40 o to 140 o C: 18 x 2.09 x 40 = 1.5 kJ (heating the ice) 1.0 x 6.01 = 6.0 kJ (melting the ice) 18 x 4.18 x 100 = 7.5 kJ (heating the water) 1.0 x 40.7 = 40.7 kJ (boiling the water) 18 x 1.84 x 40 = 1.3 kJ (heating the steam)

33 Properties of Liquids Viscosity Viscosity is the resistance of a liquid to flowing. High viscosity liquids (e.g. molasses, motor oil) flow slowly. Low viscosity liquids (e.g. water, gasoline) flow quickly.

34 Surface tension is a measure of the strength of intermolecular attractions which pull on molecules at the surface of a liquid. It is because of its high surface tension that water tends to “bead” up on a waxy surface. That’s because a sphere gives the minimum ratio of surface area to volume. The high surface tension of water also causes molecules on the surface to pack very closely together. This is why some insects can “walk on water”. Properties of Liquids Surface Tension

35 Structure of Solids Crystalline Solid Ordered arrangement of atoms (or molecules) in 3-dimensional structure, called a lattice e.g. Quartz (SiO 2 ) Amorphous Solid Irregular (disordered) arrangement of atoms (or molecules) e.g. Silica Glass (SiO 2 )

36 The Crystalline Lattice Lattice: Three Dimensional array of points representing the centers of each atom in the crystal. Lattice Point: Point corresponding to atom center Unit Cell: Parallelopiped corresponding to minimum unit which can be used to replicate crystal

37 Bonding in Solids There are four classifications of solids, depending on the type of bonds that are present. Covalent-Network Solids Ionic Solids Metallic Solids Molecular Solids

38 Covalent-Network Solids Form of particles: Atoms connected in network of covalent bonds Forces between particles: Covalent bonds Properties: Very Hard Very high melting point Usually poor thermal and electrical conductivity Examples: Diamond (C), Quartz (SiO 2 ) Diamond Each carbon is connected to 4 others by a covalent bond

39 Ionic Solids Form of particles: Positive and negative ions Forces between particles: Electrostatic attractions Properties: Hard and Brittle High melting point Poor thermal and electrical conductivity Examples: All typical salts. e.g. NaCl, Ca(NO 3 ) 3, MgBr

40 Metallic Solids Form of particles: Atoms Forces between particles: Metallic Bonds (due to delocalized valence electrons) Properties: Soft to very hard Low to very high melting point Excellent thermal and electrical conductivity Malleable and Ductile Examples: All metals. e.g. Cu, Fe, Sn, Au, Ag Bonding due to delocalized valence electrons (shown in blue) Strength of bonding varies between different metals, resulting in wide range of physical properties

41 Molecular Solids Form of particles: Atoms or molecules Forces between particles: Dispersion Dipole-Dipole (if molecules are polar) Hydrogen Bonds (if O-H, N-H, F-H) Properties: Fairly soft Moderately low melting point (usually <200 o C) Poor thermal and electrical conductivity Examples: Argon, CH 4, CO 2, C 6 H 12 O 6 (sucrose), H 2 O