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Intermolecular Forces

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  1. 11 Intermolecular Forces 11.1 Polarity of Molecules 11.2 Van der Waals’ Forces 11.3 Van der Waals’ Radii 11.4 Molecular Crystals 11.5 Hydrogen Bonding

  2. 11.1 Polarity of Molecules

  3. electrostatic attraction between dipoles, i.e. the attraction between the +ve end of one molecule and the -ve end of another molecule 11.1 Polarity of molecules (SB p.275) Polarity of molecules (very weak when compared with covalent bond between atoms in molecule) Intermolecular forces Van der Waal’s forces hydrogen bonding

  4. 11.1 Polarity of molecules (SB p.275) Polarity of molecules 3 types of dipoles Permanent dipole Instantaneous dipole Induced dipole

  5. 11.1 Polarity of molecules (SB p.275) Permanent dipole A permanent dipole exists in all polar molecules as a result of the difference in the electronegativity of bonded atoms.

  6. 11.1 Polarity of molecules (SB p.276) Instantaneous dipole An instantaneous dipole is a temporary dipole that exists as a result of fluctuation in the electron cloud.

  7. 11.1 Polarity of molecules (SB p.276) Induced dipole An induced dipole is a temporary dipole that is created due to the influence of neighbouring dipole (which may be a permanent or an instantaneous dipole).

  8. 11.2 Van der Waals’ Forces

  9. 11.2 Van der Waals’ forces (SB p.276) Van der Waals’ Forces Van der Waals’ forces Dipole-Dipole Interaction Dipole-Induced Dipole Interaction Instantaneous Dipole-Induced DipoleInteraction

  10. 11.2 Van der Waals’ forces (SB p.277) Dipole-dipole interactions • Polar molecules have permanent dipole moments. • They tend to orient themselves in such a way that the attractive forces between molecules are maximized while repulsive forces are minimized

  11. 11.2 Van der Waals’ forces (SB p.277) Dipole-induced dipole interactions • When a non-polar molecule approaches a polar molecule (with a permanent dipole), a dipole will be induced in the non-polar molecule • The dipole induced will be in opposite orientation to that of the polar molecule.

  12. 11.2 Van der Waals’ forces (SB p.277) Instantaneous dipole-induced dipole interactions • The instantaneous dipole will induce a dipole moment in the neighbouring atom by attracting opposite charges • If the +ve end of the dipole is pointing towards a neighbouring atom, the induced dipole will then have its -ve end pointing towards the +ve pole of that dipole

  13. 11.2 Van der Waals’ forces (SB p.278) Instantaneous dipole-induced dipole interactions

  14. 11.2 Van der Waals’ forces (SB p.279) Strength of van der Waals’ forces

  15. 11.2 Van der Waals’ forces (SB p.279) Strength of van der Waals’ forces • Two factors affecting the strength of van der Waals’ forces • Sizes of electron clouds of molecules • Surface area of molecules

  16. 11.2 Van der Waals’ forces (SB p.279) 1. Size of electron cloud The greater the no. of e-s in a molecule The more weakly they are held by the nucleus The easier the instantaneous dipole can be set up (greater van der Waals’ forces)

  17. 11.2 Van der Waals’ forces (SB p.280) 2. Surface area of molecule

  18. 11.2 Van der Waals’ forces (SB p.280) 2. Surface area of molecule The van der Waals’ forces also increase with the surface area of the molecule.

  19. Let's Think 1 11.2 Van der Waals’ forces (SB p.280) Change of states and intermolecular forces • 3 different states: solid, liquid and gas • Molecules in different orders in the three states •  highest order in the solid state •  lowest order in the gas state • Change of state is related to the strength of intermolecular forces of the molecular substances

  20. Line TC: Boiling point curve Line TB: Melting point curve Critical point: Beyond this pt, liquid and vapour become indisguishable Triple point: 3 states coexist at equilibrium 11.2 Van der Waals’ forces (SB p.281) Pressure-temperature diagram of carbon dioxide

  21. Check Point 11-2 11.2 Van der Waals’ forces (SB p.280) Distinguishing features of the pressure-temperature diagram of carbon dioxide • Melting point curve has a +ve slope •  melting of CO2 becomes more difficult with increase in temp. • Triple point is at 5.1 atm and –57 oC •  at 1 atm, CO2 sublimes •  No liquid state of CO2 exists under normal atmospheric condition •  Dry ice

  22. 11.3 Van der Waals’ Radii

  23. 11.3 Van der Waals’ radii (SB p.282) Van der Waals’ Radii • Van der Waals’ forces determine the closest distance between argon atoms

  24. Covalent and van der Waals’ radii of iodine 11.3 Van der Waals’ radii (SB p.283) Radii of iodine • The covalent radius is one half of the distance between two atoms in the same molecule. • The van der Waals’ radius is one half of the distance between two atoms in adjacent molecule.

  25. 11.3 Van der Waals’ radii (SB p.283) Covalent and van der Waals’ radii of some elements

  26. Check Point 11-3 Sum of covalent radii of two C atoms Sum of van der Waals’ radii of two C atoms 11.3 Van der Waals’ radii (SB p.284) Structure of graphite

  27. 11.4 Molecular Crystals

  28. 11.4 Molecular crystals (SB p.284) Molecular crystals A molecular crystal is a structure which consists of individual molecules packed together in a regular arrangement by weak intermolecular forces.

  29. 11.4 Molecular crystals (SB p.285) Iodine A unit cell of iodine crystal showing the orientation of I2 molecules

  30. 11.4 Molecular crystals (SB p.285) Dry ice A unit cell of dry ice (CO2)

  31. 11.4 Molecular crystals (SB p.286) Buckminsterfullerene • New form carbon, C60

  32. 11.5 Hydrogen Bonding

  33. F being very electronegative very +ve F atom being small enough to approach very close to the H atom in the neighbouring molecule 11.5 Hydrogen bonding (SB p.286) HF molecule

  34. 11.5 Hydrogen bonding (SB p.286) Relative strength of van der Waals’ forces, hydrogen bond and covalent bond

  35. 11.5 Hydrogen bonding (SB p.287) Formation of hydrogen bonds in hydrogen fluoride

  36. 11.5 Hydrogen bonding (SB p.287) Formation of hydrogen bonds in water

  37. 11.5 Hydrogen bonding (SB p.287) Formation of hydrogen bonds in ammonia

  38. 11.5 Hydrogen bonding (SB p.287) Formation of hydrogen bonds in methanol

  39. 11.5 Hydrogen bonding (SB p.287) Essential requirements for the formation of hydrogen bond • H atom must be directly bonded to a highly electronegative atom (e.g. F, O and N) • An unshared pair of electrons (lone pair electrons) is present on the electronegative atom

  40. 11.5 Hydrogen bonding (SB p.288) Pressure-temperature diagram of water • Quite similar to that of CO2 • One exception: slope of melting point curve is negative

  41. 11.5 Hydrogen bonding (SB p.288) Extraordinary features in relation to hydrogen bond formation • High m.p. and b.p. • Ice melts to give liquid water with a contraction in volume

  42. Abnormally high b.p. of NH3, H2O and HF 11.5 Hydrogen bonding (SB p.288) Importance of hydrogen bonding in physical phenomena 1. Anomalous properties of the second period hydrides

  43. 11.5 Hydrogen bonding (SB p.288) • Explanation: • High electronegativities of F(4.0), N(3.0) and O(3.5) • Formation of intermolecular hydrogen bonds • Hydrogen bonds are much stronger than van der Waals’ forces •  more energy is needed to break the hydrogen bonds •  abnormally high b.p.

  44. 11.5 Hydrogen bonding (SB p.289) Enthalpy of vaporization • Energy required to vaporize 1 mole of liquid • Related to the strength of intermolecular forces that exist in the liquid

  45. 11.5 Hydrogen bonding (SB p.289) Enthalpy of vaporization of Group VI hydrides • Abnormally high enthalpy of vaporization •  formation of intermolecular hydrogen bonds

  46. 11.5 Hydrogen bonding (SB p.290) Boiling points and solubilities of alcohols • B.p. of thiols are much lower than those of alcohols •  formation of intermolecular hydrogen bonds between alcohol molecules

  47. Example 11-5A 11.5 Hydrogen bonding (SB p.290) Dimerization of carboxylic acids • When ethanoic acid is dissolved in non-polar solvents, the molecular mass of found to be 120 (not 60) • Formation of dimer

  48. Hydrogen bonding in water 11.5 Hydrogen bonding (SB p.291) Hydrogen bonding in water and ice • In water, the molecules are in constant motion. H bonds areformed and broken continually. The arrangement of molecules are thus in random.

  49. Let's Think 2 Hydrogen bonding in ice 11.5 Hydrogen bonding (SB p.292) • In ice, the molecular motion is of a minimum and the molecules are oriented in such a way that the max. no. of H bonds areformed. This creates an open structure. (density of ice < density of water)

  50. 11.5 Hydrogen bonding (SB p.293) Hydrogen bonding in proteins • Primary structure of protein: polymer of amino acids