1. Unit 2: Chemical Bonding Chemistry2202 1

2. Outline  Bohr diagrams  Lewis Diagrams  Types of Bonding  Ionic bonding  Covalent bonding (Molecular)  Metallic bonding  Network covalent bonding 2

3.  Types of Bonding (cont’d)  London Dispersion forces  Dipole-Dipole forces  Hydrogen Bonding  VSEPR Theory (Shapes)  Physical Properties 3

4. Bohr Diagrams (Review) How do we draw a Bohr Diagram for - The F atom? - The F ion? Draw Bohr diagrams for the atom and the ion for the following: AlSC lBe 4

5. Lewis Diagrams LD provide a method for keeping track of electrons in atoms, ions, or molecules Also called Electron Dot diagrams the nucleus (p+ & n 0 ) and filled energy levels are represented by the element symbol 5

6. Lewis Diagrams dots are placed around the element symbol to represent valence electrons 6

7. Lewis Diagrams eg. Lewis Diagram for F F lone pair bonding electron 7

8. Lewis Diagrams lone pair – a pair of electrons not available for bonding bonding electron – a single electron that may be shared with another atom 8

9. Lewis Diagrams eg. Lewis Diagram for C C 9

10. Lewis Diagrams eg. Lewis Diagram for P P 10

11. Lewis Diagrams eg. Lewis Diagram for Na Na 11

12. Lewis Diagrams For each atom draw the Lewis diagram and state the number of lone pairs and number of bonding electrons LiBeAlSi MgNBO 12

13. Lewis Diagrams for Compounds draw the LD for each atom in the compound The atom with the most bonding electrons is the central atom Connect the other atoms using single bonds (1 pair of shared electrons) In some cases there may be double bonds or triple bonds 13

14. Lewis Diagrams for Compounds eg. Draw the LD for: PH 3 CF 4 Cl 2 O C2H6C2H6 C2H4C2H4 C2H2C2H2 14

15. Lewis Diagrams for Compounds eg. Draw the LD for: NH 3 SiCl 4 N 2 H 4 HCN SI 2 CO 2 N 2 H 2 CH 2 O POICH 3 OH N 2 H 2 O 2 15

16. Lewis Diagrams for Compounds A structural formula shows how the atoms are connected in a molecule. To draw a structural formula:  replace the bonded pairs of electrons with short lines  omit the lone pairs of electrons 16

17. Why is propane (C 3 H 8 ) a gas at STP while kerosene (C 10 H 22 ) a liquid? 17

18. Why is graphite soft enough to write with while diamond is the hardest substance known even though both substances are made of pure carbon? 18

19. Why can you tell if it is ‘real gold’ or just ‘fool’s gold’ (pyrite) by hitting it with a rock? 19

20. ‘As Slow As Cold Molasses’ ‘All Because of Bonding’ 20

21. ‘liquids’ @ -30 ºC 21

22. Bonding Bonding between atoms, ions and molecules determines the physical and chemical properties of substances. Bonding can be divided into two categories: - Intramolecular forces - Intermolecular forces 22

23. Bonding Intramolecular forces are forces of attraction between atoms or ions. Intramolecular forces include: 1. ionic bonding 2. covalent bonding 3. metallic bonding 4. network covalent bonding 23

24. Bonding Intermolecular forces are forces of attraction between molecules. Intermolecular forces include: 5. London Dispersion Forces 6. Dipole-Dipole forces 7. Hydrogen Bonding 24

25. Ionic and Covalent Bonding ThoughtLab p. 161 Identify #’s 1 - 6 25

26. Ionic Bonding Occurs between cations and anions – usually metals and non-metals. An ionic bond is the force of attraction between positive and negative ions. Properties:  conduct electricity as liquids and in solution  hard crystalline solids  high melting points and boiling points  brittle 26

27. In an ionic crystal the ions pack tightly together. The repeating 3-D distribution of cations and anions is called an ionic crystal lattice. Ionic Bonding 27

28. Each anion can be attracted to six or more cations at once. The same is true for the individual cations. Ionic Bonding 28

29. Ionic Bonding 29

30. Covalent Bonding Occurs between non-metals in molecular compounds. Atoms share bonding electrons to become more stable (noble gas structure). A covalent bond is a simultaneous attraction by two atoms for a common pair of valence electrons. 30

31. Covalent Bonding Molecular compounds have low melting and boiling points. Exist as distinct molecules. 31

32. Covalent Bonding Molecular compounds do not conduct electric current in any form 32

33. PropertyIonicMolecular Type of elements Force of Attraction Electron movement State at room temperature Metals and nonmetals Non-Metals Positive ions attract negative ions Atoms attract a shared electron pair Electrons move from the metal to the nonmetal Electrons are shared between atoms Always solids Solids, liquids, or gas 33

34. PropertyIonicMolecular Solubility Conductivity in solid state Conductivity in liquid state Conductivity in solution Soluble or low solubility Soluble or insoluble None Conducts 34

35. Metallic Bonding (p. 171) metals tend to lose valence electrons. valence electrons are loosely held and frequently lost from metal atoms. This results in metal ions surrounded by freely moving valence electrons. metallic bonding is the force of attraction between the positive metal ions and the mobile or delocalised valence electrons 35

36. Metallic Bonding 36

37. Metallic Bonding This theory of metallic bonding is called the ‘Sea of Electrons’ Model or ‘Free Electron’ Model 37

38. Metallic Bonding This theory accounts for properties of metals 1. electrical conductivity - electric current is the flow of electrons - metals are the only solids in which electrons are free to move 2. solids - Attractive forces between positive cations and negative electrons are very strong 38

39. Metallic Bonding 3. malleability and ductility - metals can be hammered into thin sheets(malleable) or drawn into thin wires(ductile). - metallic bonding is non-directional such that layers of metal atoms slide past each other under pressure. 39

40. Network Covalent Bonding (p. 199) occurs in 3 compounds (memorize these)  diamond – C n  carborundum – SiC  quartz – SiO 2 large molecules with covalent bonding in 3-d each atom is held in place in 3-d by a network of other atoms 40

41. Network Covalent bonding Properties:  the highest melting and boiling points  the hardest substances  brittle  do not conduct electric current in any form 41

42. Strongest 1. Network Covalent (C n,SiO 2, SiC) 2. Ionic bonding(metal & nonmetal) 3. Metallic bonding (metals) 4. Molecular (nonmetals) Weakest MP & BP decreases 42

43. Valence Shell Electron Pair Repulsion theory (VSEPR) The shape of molecules is caused by repulsion between valence electron pairs around the atoms in a compound. Repulsion between valence electron pairs force them to move as far away from each other as possible. 43

44. Valence Shell Electron Pair Repulsion theory (VSEPR) To determine molecular shapes, count the # of bonds and # of lone pairs on the central atom(s). We will examine 5 molecular shapes 44

45. 1. Tetrahedral (4 bonds; 0 lone pairs) 45

46. 2. Pyramidal (3 bonds; 1 lone pair) 46

47. 3. V-shaped (2 bonds; 2 lone pairs) 47

48. 4. Trigonal Planar (3 bonds; 0 lone pairs) 48

49. 5. Linear (2 bonds; 0 lone pairs) 49

50. For each molecule below draw the Lewis diagram and the shape diagram. HOClH 2 SeH 2 O 2 NBr 3 C 2 F 4 C 2 H 6 CHCl 3 CH 3 OH PBr 3 I 2 SiH 4 HCN SiH 2 OC 2 H 2 50

51. Electronegativity - EN - p. 174 EN measures the attraction that an atom has for shared electrons. A higher EN means a stronger attraction or electrostatic pull on valence electrons EN values increase as you move: - from left to right in a period - up in a group or family 51 Electronegativity & Covalent Bonds

52. Increases 52

53. Electronegativity & Covalent Bonds polar covalent bond - a bond between atoms with different EN - the bonding pair is attracted more strongly to the atom with the higher EN ClH δ− δ+ 53 bond dipole

54. nonpolar covalent bond - a bond between atoms with the same EN - the bonding pair is shared equally between the atoms Complete: #’s 7 – 9 on p.178 54 Electronegativity & Covalent Bonds

55. polar molecule - a molecule in which the bond dipoles do not cancel each other - a polar molecule has a molecular dipole that points toward the more electronegative end of the molecule. eg. HCN 55

56. Electronegativity & Covalent Bonds nonpolar molecule - a molecule in which the bond dipoles cancel each other OR - there are no bond dipoles eg. CO 2 PH 3 56

57. Electronegativity & Covalent Bonds To determine whether a molecule is polar: - draw the Lewis diagram & shape diagram - draw the bond dipoles & determine whether they cancel 57

58. Intermolecular Forces 58

59. Strongest 1. Network Covalent (C n,SiO 2, SiC) 2. Ionic bonding(metal & nonmetal) 3. Metallic bonding (metals) 4. Molecular (nonmetals) Weakest MP & BP decreases 59

60. To compare mp and bp in covalent compounds you must use: - London Dispersion forces (p. 204) (in all molecules) - Dipole-Dipole forces (pp. 202, 203) (in polar molecules) - Hydrogen Bonding (pp. 205, 206) (when H is bonded to N, O, or F) 60

61. Intermolecular Forces (p. 202) 61

62. Intermolecular Forces Covalent compounds have low mp and bp because attractive forces between molecules are very weak. Intermolecular forces were studied extensively by the Dutch physicist Johannes van der Waals In his honor, two types of intermolecular force are called Van der Waals forces. 62

63. Intermolecular Forces Intermolecular forces can be used to account for the physical properties of covalent compounds. 63

64. Intermolecular Forces 64 FFFF

65. 1. London Dispersion Forces LD forces exist in ALL molecular elements & compounds. The positive charges in one molecule attract the negative charges in a second molecule. The temporary dipoles caused by electron movement in one molecule attract the temporary dipoles of another molecule. 65

66. 1. London Dispersion Forces The strength of these forces depends on: a)the number of electrons more electrons produce stronger LD forces that result in higher mp and bp eg. CH 4 is a gas at room temperature. C 8 H 18 is a liquid at room temperature. C 25 H 52 is a solid at room temperature. Account for the difference. 66

67. 1. London Dispersion Forces Two molecules that have the same number of electrons are isoelectronic eg. C2H6 C2H6 and CH 3 F 67

68. 1. London Dispersion Forces b)shape of the molecule molecules that “fit together” better will experience stronger LD forces eg. Cl 2 vaporizes at -35 ºC while C 4 H 10 vaporizes at -1 ºC. Use bonding to account for the difference. 68

69. 2. Dipole-dipole Forces - occur between polar molecules - the δ+ end of one polar molecule is attracted to the δ- end of another polar molecule (& vice-versa) eg. Which has the higher boiling point CH 3 F or C 2 H 6 ? 69

70. p. 202 70

71. 3. Hydrogen Bonds - a special type of dipole-dipole force (about 10 times stronger) - only occurs between molecules that contain a H atom which is directly bonded to F, O, or N ie. the molecule contains at least one H-F, H-O, or H-N covalent bond. 71

72. 3. Hydrogen Bonds -the hydrogen bond occurs between the H atom of one molecule and the N, O, or F of a second molecule. eg. Arrange these from highest to lowest boiling point C 3 H 8 C 2 H 5 OHC 2 H 5 F 72

73. p. 206 73

74. Bonding – last worksheet 1. Use intermolecular forces to explain the following: a) Ar boils at -186 °C and F 2 boils at -188 °C. b) Kr boils at -152 °C and HBr boils at -67 °C. c) Cl 2 boils at -35 °C and C 2 H 5 Cl boils at 13 °C. 2. Examine the graph on p. 210: a) Account for the increase in boiling point for the hydrogen compounds of the Group IV elements. b) Why is the trend different for the hydrogen compounds of the Group V, VI, and VII elements? c) Why are the boiling points of the Group IVA compounds consistently lower than the others. 74

75. p. 210 75

76. 3.Which substance in each pair has the higher boiling point. Justify your answers. (a)SiC or KCl (b)RbBr or C 6 H 12 O 6 (c)C 3 H 8 or C 2 H 5 OH (d)C 4 H 10 or C 2 H 5 Cl 76

77. p. 226 #’s 13 & 14 77

78. Bond Energy (pp. 179-180) 1. Describe the forces of attraction and repulsion present in all bonds. 2. What is bond length? 3. Define bond energy. 4. Which type of bond has the most energy? 5. How can bond energy be used to predict whether a reaction is endothermic or exothermic? 78

79. Test Outline Bohr Diagrams (atoms & ions) Lewis Diagrams (Electron Dot) Ion Formation Ionic Bonding, Structures & Properties Covalent Bonding, Structures & Properties 79

80. Test Outline Metallic Bonding Theory& Properties Network Covalent Bonding & Properties Electronegativity Bond Dipoles & Polar Molecules VSEPR Theory LD, DD, & H-bonding Predicting properties (bp, mp, etc.) 80