1. Intermolecular Attractions and the Properties of Liquids and Solids

2. 2 Chapter 12 Intermolecular Forces Important differences between gases, solids, and liquids: – Gases Expand to fill their container – Liquids Retain volume, but not shape – Solids Retain volume and shape At room temperature, some are solid, others are liquid, others are gaseous. Why?

3. Physical Properties of Gases, Liquids and Solids determined by – How tightly molecules are packed together – Strength of attractions between molecules

4. 4 Inter vs. Intra-Molecular Forces Intramolecular forces – Covalent bonds within molecule – Strong –  H bond (HCl) = 431 kJ/mol Intermolecular forces – Attraction forces between molecules – Weak –  H vaporization (HCl) = 16 kJ/mol Covalent Bond (strong) Intermolecular attraction (weak)

5. When substance melts or boils – Intermolecular forces are broken – Not covalent bonds Responsible for existence of condensed states of matter Responsible for bulk properties of matter – Boiling Points and Melting Points

6. 6 Electronegativity Review Electronegativity: Measure of attractive force that one atom in a covalent bond has for electrons of the bond

7. 7 Bond Dipoles Two atoms with different electronegativity values share electrons unequally Electron density is uneven – Higher charge concentration around more electronegative atom Bond dipoles – Indicated with delta (δ) notation – Indicates partial charge has arisen

8. 8 Three Important Types of Intermolecular Forces 1. Dipole-dipole forces – Hydrogen bonds 2. London dispersion forces 3. Ion-dipole forces – Ion-induced dipole forces

9. 9 I. Dipole-dipole Attractions Occur only between polar molecules – Possess dipole moments Molecules need to be close together Polar molecules tend to align their partial charges – + to – As dipole moment , intermolecular force  +   + + 

10. 10 I. Dipole-dipole Attractions Tumbling molecules – Mixture of attractive and repulsive dipole-dipole forces – Attractions (- -) greater than repulsions(- -) – Get net attraction – ~ 1% of covalent bond

11. 11 Hydrogen Bonds Special type of Dipole-Dipole Interaction – Very strong dipole-dipole attraction – ~40 kJ/mol Occurs between H and highly electronegative atom (O, N, or F) – H—F, H—O, and H—N bonds very polar Positive end of one can get very close to negative end of another

12. 12 Examples of Hydrogen Bonding

13. 13 Effects of Hydrogen Bonding Boiling points of H compounds of elements of Groups IVA, VA, VIA, and VIIA. Boiling points of molecules with H bonding are higher than expected. Don’t follow rule that BP  as MM  (London forces  ) Boiling Point (°C)

14. 14 Hydrogen Bonding in Water Responsible for expansion of water as it freezes Hydrogen bonding produces strong attractions in liquid Hydrogen bonding (dotted lines) between water molecules in ice form tetrahedral configuration

15. 15 II. London Dispersion Forces Intermolecular forces between nonpolar molecules Two neutral molecules (atoms) can affect each other – Nucleus of 1 molecule (atom) attracts e  ’s of adjacent molecule (atom) – Electron cloud distorts – Temporary or instantaneous dipole forms – One instantaneous dipole can induce another in adjacent molecule (atom) – Results in net attractive force

16. 16 London Dispersion Forces Instantaneous dipole-induced dipole attractions – London Dispersion Forces – London forces – Dispersion forces Decrease as 1/d 6 (d = distance between molecules) Effect enhanced with increased particle mass Operate between all molecules – Neutral or net charged – Nonpolar or polar

17.  London Forces  as MM   More e , less tightly held  London Forces  as electron cloud volume (size)  Larger molecules have stronger London forces and thus higher boiling points.

18. 18 2. Number of Atoms in Molecule London forces depend on number atoms in molecule Boiling point of hydrocarbons demonstrates this trend Formula BP at 1 atm,  C Formula BP at 1 atm,  C CH 4  161.5 C 5 H 12 36.1 C2H6C2H6  88.6 C 6 H 14 68.7 C3H8C3H8  42.1 :: C 4 H 10  0.5 C 22 H 46 327

19. 19 III. Ion-dipole Attractions Attractions between ion and charged end of polar molecules – Attractions can be quite strong as ions have full charges (a) Negative ends of water dipoles surround cation (b) Positive ends of water dipoles surround anion

20. 20 Ex. Ion-dipole Attractions AlCl 3 ·6H 2 O  Positive charge of Al 3+ ion attracts partial negative charges  – on O of water molecules  Ion-dipole attractions hold water molecules to metal ion in hydrate  Water molecules are found at vertices of octahedron around aluminum ion Attractions between ion and polar molecules

21. 21 Using Intermolecular Forces Often can predict physical properties (like BP and MP) by comparing strengths of intermolecular attractions – Ion-Dipole – Hydrogen Bonding – Dipole-Dipole – London Dispersion Forces Larger, longer, heavier molecules have stronger IMFs Smaller, more compact, lighter molecules have weaker IMFs Weakest Strongest

22. 22 Phase Changes Changes of physical state – Deal with motion of molecules As temperature changes – Matter will undergo phase changes Liquid  Gas – Evaporation – As heat H 2 O, forms steam or water vapor – Requires energy or source of heat to occur

23. 23 Phase Changes Solid  Gas – Sublimation – Ice cubes in freezer, leave in long enough disappear – Endothermic Gas  Liquid – Cooling or Condensation – Dew is H 2 O vapor condensing onto cooler ground – Exothermic

24. 24 Phase Changes Energy of System Gas Solid Liquid Melting or Fusion VaporizationCondensation Freezing Sublimation Deposition  Exothermic, releases heat  Endothermic, absorbs heat

27. 27 Rate of Evaporation Depends on – Temperature – Surface area – Strength of intermolecular attractions Molecules that escape from liquid have larger than average KE’s When they leave – Average KE of remaining molecules is less – T lower

28. 28 Effect of Temperature on Evaporation Rate For given liquid – Rate of evaporation per unit surface area  as T  Why? – At higher T, total fraction of molecules with KE large enough to escape is larger – Result: rate of evaporation is larger

29. 29 Kinetic Energy Distribution in Two Different Liquids Smaller IMF’s Lower KE required to escape liquid A evaporates faster Larger IMF’s Higher KE required to escape liquid B evaporates slower AB

30. 30 Vapor Pressure Diagram

32. T-t curves

33. Supercooling

34. 34 Phase Diagram of Water