Lecture outline: Chapter 11 Intermolecular attractive forces

Transcription

1 Lecture outline: Chapter 11 Intermolecular attractive forces Intermolecular forces Phase changes Vapor pressure Phase diagrams Types of solids 1

2 Gases The volume occupied by gas molecules is much less than the volume in which they reside Molecules of gases are in continual random motion Attractive/repulsive forces between gas molecules are negligible The average kinetic energy of a gas molecule is proportional to the temperature Which of these properties apply to liquids and solids as well?? 2

3 Physical state Defined volume? Defined shape? Compress/ expand? Gas Liquid Solid 3

4 Attractive forces Intramolecular forces: attractive forces of atoms within compounds (think: rope knot, zipper) Molecular compounds: covalent bonds, bond energies Ionic compounds: ionic bonds, lattice energies Intermolecular forces: attractive forces between different molecules (think: velcro, snap button) δ H δ+ Cl δ H δ+ Cl All attractive forces are electrostatic in nature 4

5 Magnitudes of attractive forces Bond energies Lattice energies Intermolecular attractions 5

6 Intermolecular forces Iondipole Dipoledipole London dispersion Hydrogen bonding Attraction between an ion and a molecule Attraction between molecules All attractive forces are electrostatic in nature 6

7 Review electronegativity, polarity, and dipole moment Increases from left to right in a period 7

8 Review electronegativity, polarity, and dipole moments δ + δ H F Polar less EN more EN Net dipole δ O δ + δ C O No net dipole Nonpolar Net dipole δ O δ + H H δ + Polar 8

9 Iondipole force The attraction between an ion and the partial charge on an end of a polar neutral molecule 9

10 IonDipole Forces δ + δ δ + δ + δ δ + δ δ + 10

11 IonDipole Forces δ δ + δ δ + δ + δ δ + δ 11

12 IonDipole Forces involving water δ + Cl δ + δ + Cl δ δ δ + δ + δ Na + Cl δ + δ δ + Cl δ + 12

13 Dipoledipole forces Attraction between neutral, polar molecules δ H δ+ Cl δ H δ+ Cl 13

14 DipoleDipole Forces + H Cl Cl Cl + H Cl + H + H Cl Cl H + H + Cl Cl H + + H Cl H + 14

15 London Dispersion forces Attraction between neutral, nonpolar molecules (or individual atoms) Since all attractions are electrostatic, how is this possible?? δ 0 δ 0?? δ 0 δ 0 15

16 London Dispersion forces Attraction between neutral, nonpolar molecules (or individual atoms) The motion of electrons in atoms and molecules can result in a shortlived instantaneous dipole e density shifts to left δ δ + δ 0 δ 0 δ 0 δ 0 e density shifts to right δ + δ 16

17 London Dispersion forces The motion of electrons in atoms and molecules can result in a shortlived instantaneous dipole An instantaneous dipole on one molecule (or atom) can induce an instantaneous dipole on an adjacent molecule (or atom) resulting in the dipoledipole attraction δ + 0 δ 0 δ + 0 δ 0 δ + 0 δ 0 δ + 0 δ 0 17

18 London Dispersion forces The motion of electrons in atoms and molecules can result in a shortlived instantaneous dipole An instantaneous dipole on one molecule (or atom) can induce an instantaneous dipole on an adjacent molecule (or atom) resulting in the dipoledipole attraction A switch in the direction of the instantaneous dipole on one molecule results in a switch in the direction on adjacent molecules δ + δ + δ δ + δ + δ + δ + δ + δ + 18

19 London Dispersion forces δ + δ δ + δ δ + δ δ + δ δ δ + δ δ + δ δ + δ δ + δ + δ δ + δ δ + δ δ + δ 19

20 London Dispersion Forces can also operate in a collection of neutral atoms e density shifts to left δ δ + δ 0 δ0 e density shifts to right δ + δ δ δ + δ δ + δ δ + δ δ + δ δ + δ + δ δ + δ δ + δ δ + δ δ + δ 20

21 Magnitude of London dispersion forces and size.. 21

22 Boiling and melting points are indicators of the strength of an intermolecular attractive force Δ melting Δ boiling 22

23 Boiling points of the Noble gases: He 4.2 K Ne 27.1 K Ar 87.3 K Kr K Xe Rn K 23

24 Hydrocarbon boiling points H A. B. C. H H C H H C C H H C C H H H H C H H H bp = 161 C bp = 89 C bp = 44 C H H H H D. H H H H E. H H H H H H C C C C H H C C C C C H H H H H H H H H H bp = 0.5 C bp = 36 C 24

25 Effect of shape on hydrocarbon boiling points: molecules with formula C 5 H 12 bp = 36.0 C bp = 27.7 C bp = 9.5 C 25

26 Hydrogen bonding Special type of dipoledipole force The intermolecular attraction between a H atom in a very polar bond and an unshared e pair on a small electronegative atom The intermolecular attraction between a H atom attached to N, O, or F and a N, O, or F atom 26

27 Hydrogen bonding in water δ + δ δ + δ + δ + δ δ + δ δ δ + δ + δ δ + δ + 27

28 Examples of hydrogen bonds Between waters Between ammonias Between water and ammonia Between water and ethylene glycol 28

29 Hydrogen bonding in ice 29

30 Hydrogen bonding in ice ICEShex saemod2.pse 30

31 Hydrogen bonding in ice 31

32 Hydrogen bonding in ice ICEShex saemod2.pse 32

33 Hydrogen bonding in liquid water At 25 C, thermal energy closely matches H bond energy (20 kj/mol) Average lifetime of aqueous Hbond is s Disordered 3dimensional network Each water has an average of 3.4 Hbonds 33

34 More than one type of force can contribute to intermolecular attractions in a molecule Consider: (1) size of molecule (2) shape of molecule Amount of accessible valence electron density to overlap another molecule (3) polarity of molecule ( EN) 34

35 boiling point ( C) More than one type of force can contribute to intermolecular attractions in a molecule Boiling points of group 6A hydrides H2O H2Te H2S H2Se Molar mass (g/mol) Consider: (1) size of molecule (2) shape of molecule (3) polarity of molecule ( EN) 35

36 More than one type of force can contribute to intermolecular attractions in a molecule Boiling points of hydrogen halides HF Consider: (1) size of molecule (2) shape of molecule boiling point ( C) HI (3) polarity of molecule ( EN) HBr HCl Molar mass (g/mol) 36

37 Intra and intermolecular attractive force strengths Remember it always takes energy to break a bond, whether it is intra or intermolecular covalent bond (bond E): ionic bond (lattice E): iondipole Hydrogen bond dipoledipole London dispersion ~ kj/mol ~ kj/mol ~15 kj/mol ~525 kj/mol ~210 kj/mol ~010 kj/mol 37

38 Intra and intermolecular forces in water H δ + H δ+ 463 kj/mol δ O δ O δ + H H δ + ~19 kj/mol 38

39 Intra and intermolecular forces in 463 kj/mol water ~19 kj/mol 39

40 Intermolecular attractions in neutral molecules In general, the relative strengths are: Hbond > dipoledipole > Londondispersion Strengths of dipoledipole and Londondispersion forces are related to the polarity, size, and shape of the molecule Boiling points are good indicators of the strength of an intermolecular attractive force: the higher the bp, the stronger the attraction; the lower the bp, the weaker the attraction 40

41 To visualize intermolecular forces, reference everything to absolute zero, where molecular motion is at a minimum, and everything is solid. Then think about adding heat, and ask what happens 41

42 What is the major type of intermolecular HF attractive force between molecules of: F 2 PCl 3 BrF 42

43 Which substance in the following pairs would you predict has a higher boiling point? HF or HBr CH 4 or C 4 H 10 NH 3 or PH 3 H 2 O or H 2 S MgBr 2 or PBr 3 43

44 Important properties of liquids Melting and boiling points Viscosity Surface tension Ability to form mixtures Vapor pressure 44

45 Viscosity: resistance of a liquid to flow Hydrocarbon Formula Boiling point Viscosity (cp) Hexane Heptane Octane Nonane Decane

46 Surface tension of a liquid Surface molecules are in a different environment than interior molecules 46

47 Surface tension of a liquid Author Michael Apel. T This file is licensed under the Creative Commons AttributionShare Alike 3.0 Unported license. Author: PD. This file is licensed under the Creative Commons Attribution Share Alike 3.0 Unported license. Attribution: Chris 73 / Wikimedia Commons This file is licensed under the Creative Commons AttributionShare Alike 3.0 Unported license. 47

48 Changes of state Energy of system gas liquid solid 48

49 exothermic endothermic Δ Δ 49 condensation liquid freezing gas deposition solid Energy of system sublimation melting (fusion) vaporization

50 Energy changes associated with changes of state Δ endothermic Δ endothermic exothermic Δ Δ How does evaporative cooling work? exothermic 50

51 Energy changes associated with changes of state Why does your hand get burned when exposed to gaseous water (steam) but not when exposed to hot air at the same temperature? 51

52 Heating curve for water Temperature ( C) liquid lines in red: raising temperature of (heating) a single phase gas liquid gas 0 25 ice ice liquid heat added (kj) 52

53 Heating curve for water Temperature ( C) ice liquid ice liquid heat added (kj) gas lines in green: phase transitions: all heat goes into breaking bonds, so no T increase liquid gas 53

54 The heat input for each region of the curve is determined by specific heats and enthalpies of phase transitions Temperature ( C) SH gas = 1.84 J/(g C) SH liquid = 4.18 J/(g C) liquid gas ΔH vap = kj/mol 0 25 ice ΔH fusion = 6.00 kj/mol SH ice =2.09 J/(g C) heat added (kj) 54

55 Use specific heat values, and ΔH vaporization and ΔH fusion values, to calculate ΔH for processes involving change of state (see chapter 11 self test) 55

56 States of matter are dependent on both temperature and pressure Critical temperature: the highest temperature at which a substance can exist as a liquid Critical pressure: the pressure required to liquify a substance at it s critical temperature 56

57 Which of the following substances can be liquefied at room temperature (298 K), given sufficient pressure? Substance Critical T (K) Critical P (atm) N Ar O CH CO Propane (C 3 H 8 ) Ammonia H 2 O

58 Which of the following substances can be liquefied at room temperature (298 K), given sufficient pressure? Substance Critical T (K) Critical P (atm) N Ar O CH CO Propane (C 3 H 8 ) Ammonia H 2 O Compressed gas vs. liquified gas 58

59 Vapor pressure 59

60 Distribution of kinetic energies for molecules in a liquid T 1 Kinetic energy required to escape from liquid fraction of molecules T 2 kinetic energy 60

61 Vapor pressure The pressure exerted by molecules in the gas phase above a liquid when the vapor and liquid are at equilibrium Equilibrium has been reached when rate escape = rate return 61

62 What is the relation between the strength of intermolecular forces and the value of vapor pressure at a given temperature? 62

63 What is the relation between temperature and the value of vapor pressure for a given liquid? Equilibrium at T 1 Increase T: rate escape > rate return Equilibrium reestablished at T 2 63

64 Boiling point: the temperature at which the vapor pressure of a liquid is equal to the external pressure 1000 vapor pressure (torr) Diethyl ether Ethanol Temperature ( C) Water Ethylene glycol 64

65 Relationships between altitude, atmospheric pressure, and boiling point Location Elevation (ft) P (torr) P (atm) Boiling point ( C) Summit of Mt. Everest 29, Summit of King s peak, UT 13, USU Campus Logan, UT 4, Madison, WI San Diego, CA Surface of Pacific Ocean Death Valley, CA A medical autoclave N/A

66 A phase diagram shows the relation between pressure, temperature, and physical state of a substance solid Melting (fusion) liquid Critical point freezing Pressure vaporization condensation sublimation deposition Triple point gas Temperature 66

67 Phase diagram for CO 2 (axes not to scale) 73 solid liquid Critical point Pressure (atm) 5.11 Triple point gas Temperature ( C)

68 Phase diagram for water (axes are not to scale) 218 Critical point Pressure (atm) Triple point Temperature ( C)

69 Note the differences in direction of slope for the melting point curves for CO 2 and H 2 O Phase diagram for H 2 O Phase diagram for CO Critical point 73 solid liquid Critical point Pressure Triple point Pressure (atm) 5.11 Triple point gas Temperature Temperature ( C)

70 Water is one of only a few molecular compounds for which the solid is less dense than the liquid water Most other molecular compounds 70

71 Skate sailing on Lake Mendota, Madison Wisconsin 71

72 Solids Three types of motion for molecules: Translational (diffusion) Rotational vibrational Solids are more restricted in freedom of motion than a gas or liquid Δ melting Δ boiling 72

73 Four types of solids Molecular solid: a solid composed of atoms or molecules held together by intermolecular attractions (London, dipoledipole, or Hbonds) Covalent network solid: a solid where all atoms are connected through covalent bonds Ionic solid: a solid where cations and anions interact through ionic bonds (electrostatic attractions Metallic solid: a solid consisting of metal atoms bonded through metallic bonds c Solids can also consist of polymeric units such as polyethylene and polyester, and complex biomolecules such as cellulose and lignin 73

74 Ice: a molecular solid Other examples: sucrose (C 12 H 22 O 11 ), paraffin, I 2, dry ice (CO 2 ) 74

75 The C allotropes diamond and graphite: Covalent network solids C atoms are sp 3, tetrahedral geometry C atoms are sp 2, trigonal planar geometry Other examples: silica quartz and glass (SiO 2 ) n 75

76 The C allotropes diamond and graphite: Covalent network solids 76

77 Buckyball (C 60 ) is an allotrope of C that exists in molecular form 77

78 NaCl: an ionic solid Cl Na + 78

79 Bonding in metallic solids? Compare Na and Ne Property Neon Sodium Atomic number Std state form Monatomic gas solid bp C 883 C mp C C Electrical No Yes conductor Reactivity No Readily oxidized Color Colorless Silverywhite Atomic radius 38 pm 190 pm 79

80 Metallic solids: Positively charged metal ions swimming in a sea of electrons

81 Metallic solids: Positively charged metal ions swimming in a sea of electrons

82 Properties of metals accounted for by the electron sea model Solids at room temperature Good electrical conductors high heat conductivity and heat capacity malleable and ductile Readily undergo oxidation