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Intermolecular Forces and the Physical Properties of Liquids and Solids.

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Presentation on theme: "Intermolecular Forces and the Physical Properties of Liquids and Solids."— Presentation transcript:

1 Intermolecular Forces and the Physical Properties of Liquids and Solids

2  Draw Lewis Structures for CCl 4 and CH 3 Cl.  What’s the same? What’s different?

3  Bonds are polar if electrons are shared unequally (differences in electronegativity) (e.g., CH 3 Cl)  Individual bonds can be polar but can cancel each other out to yield a nonpolar molecule (e.g., CO 2 )  Molecules with a lone pair of electrons or atoms with different electronegativities are polar (have a dipole moment) 3 Polarity

4  Draw Lewis Structures and Dipole Moments along each bond:  HCl  CCl 4  NH 3  H 2 O ● BF 3 ● CH 3 Cl ● H 2

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6  Bonds: attractive forces within molecules  Intermolecular Forces (IMF): attractive forces between molecules  Intermolecular forces are WEAKER than bonds (intramolecular forces), but have profound effects on the properties of liquids 431 kJ/mol 16 kJ/mol

7  What is the difference between bonds and intermolecular forces?  IMF affect boiling points, melting points, and solubilities  As a group, intermolecular forces are called van der Waals forces: ◦ London Dispersion ◦ Dipole-dipole ◦ Hydrogen bonds ◦ Ion-dipole

8  London dispersion forces: attractive forces that result from temporary shift of electrons in atoms or molecules; present in all molecules  Explains why certain nonpolar substances benzene, bromine, etc.) are liquids at room temperature. 8 Figure 10.5 I 2 induced dipoles

9  Only force in non-polar molecules and in unbonded atoms (e.g., CO 2, Ne)  Larger atoms or molecules have stronger London forces (and higher boiling points) ◦ Larger electron clouds are easier to deform (“squishier”, more polarizable) and tend to have more electrons  Found in mixtures or pure substances 9

10  Boiling point increases with the size of molecules because of increases in London forces with larger electron clouds 10

11  Pure substance or mixture made up of polar molecules  Opposite dipoles (charges) attract; like dipoles repel  Strength of these forces depends on the polarity of the molecules 11 dipole Figure 10.4

12  Like dissolves like (refers to polarity): polar liquids are more soluble in polar liquids ◦ It takes 2000 mL of H 2 O to dissolve 1 mL of CCl 4 ◦ It takes 50 mL of H 2 O to dissolve 1 mL of CH 2 Cl 2  Which member of each pair has the stronger intermolecular forces? CCl 4 or CHCl 3 CO 2 or SO 2 12

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14  Hydrogen bridge: special type of dipole-dipole force; attractive force between a hydrogen atom bonded to a very small, electronegative atom (F, O, N) and lone e - pair  Ice melting Ice melting14 HF

15  Hydrogen bridges are seen in HF, H 2 O, NH 3, but not CH 4 or H 2 S  Requirements for hydrogen bridging: ◦ H attached to a small, highly electronegative element in one molecule ◦ Small, highly electronegative element with one or more unshared electron pairs in the other molecule  Observed for the elements: F, O, N (rarely S and Cl – too large, not electronegative enough) 15

16  Molecules hydrogen bridge to themselves or to other molecules 16 bridge

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18  Which of the following pure substances will experience hydrogen bridging? H 2 O H 2 Se HBr HF NH 3 NF 3 DNA 18

19  Ions have full charges that are attracted to the partial charge on polar molecules (dipoles)  Explains solubility of ionic salts in polar solvents (e.g., NaCl in water) 19 Figure 10.3 NaCl In H 2 O

20  Strength of bonds in ionic compounds (already covered in Lattice Energy section!). NaCl CaF 2

21  Intermolecular forces generally increase in strength as: London < Dipole-Dipole < H-bridging < Ion-Dipole < Ion-ion forces (ionic bonding)  Summary: ◦ Nonpolar molecules: London Dispersion ◦ Polar molecules: Dipole-dipole ◦ Polar with H-N, H-O, or H-F: Hydrogen bridging ◦ Ionic compound: Ion-ion forces 21

22  For each substance below, indicate the strongest type of intermolecular force observed. H 2 O CO CH 4 NH 3 HCN CH 3 OH CO 2 CH 3 NH 2 F 2 N 2 22

23  Which member of each pair has stronger intermolecular forces (and higher boiling point)? CH 3 OH or CH 3 SH CH 4 or CH 3 CH 2 CH 3 CO or F 2 CH 3 Cl or HF CO 2 or NH 3 NH 3 or N 2 23 Boiling point curves

24  Which member of each pair has stronger intermolecular force (boiling point, heat of vaporization)? Explain your reasoning. CH 4 or CH 3 CH 3 NH 3 or NF 3 CO 2 or SO 2 O 2 or O 3 I 2 or Cl 2 24

25  Surface tension: attraction of molecules to each other on a liquid’s surface  Molecules must break IMF in order to move to the surface and increase the surface area (large IMF, high surface tension) 25

26  Properties (e.g., IMF) in between those of gases and solids  Viscosity: resistance to flow  Depends on intermolecular forces and sizes  Ethanol vs Glycerol 26 Viscosity glycerol vs ethanol

27  fusion (melting): s  lfreezing: l  s  vaporization: l  gcondensation: g  l  sublimation: s  gdeposition: g  s

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29  Phase diagram: plot of temperature vs. pressure; solid lines are phase equilibria Low Temp Boil Phase Diagram

30  Liquid evaporates, gaseous molecules exert a pressure (vapor pressure) that can be measured as shown below Figure Vaporizing liquid

31  Features of a phase diagram: ◦ Triple point: temperature and pressure at which all three phases (s, l, and g) are in equilibrium ◦ Vaporization curve: equilibrium between liquid and gas ◦ Melting curve: equilibrium between solid and liquid ◦ Normal melting point: melting point at 1 atm ◦ Normal boiling point: boiling point at 1 atm 31 Label Phase diagrams Atomic Phase diagrams

32  Critical point: liquid and gas phases are indistinguishable  Critical temperature, T c : highest temperature at which a substance can exist as a liquid (cannot be liquified) no matter how much pressure is applied  Critical pressure, T p : minimum pressure that must be applied to bring about liquefaction at the critical temperature

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34  Amorphous: random arrangement (rubber)  Ionic: ordered arrangement of ions (salt)

35  Molecular (left): covalent molecules in an ordered arrangement (sucrose, ice); intermolecular forces hold molecules together  Covalent network (right): atoms connected by covalent bonds in 3D array (quartz)

36  Cubic is the most common  Three forms of cubic crystals - simple, body-centered, face-centered 36 Cubic unit cells SC - build BCC –CsCl FCC – CaF 2

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