1. Chapter 9 (Silberberg 3 rd Edition) Models of Chemical Bonding 9.1 Atomic Properties and Chemical Bonds 9.2 The Ionic Bonding Model 9.3 The Covalent Bonding Model 9.4 Between the Extremes: Electronegativity and Bond Polarity 9.5 An Introduction to Metallic Bonding

2. Types of Chemical Bonding 1. What’s a Chemical Bond? Attraction that holds atoms or ions together in compounds 2. Ionic Bonding vs Covalent Bonding Ionic Bonding vs Covalent Bonding What’s the difference? Kinds of atoms involved? 3. Metallic Bonding Metallic Bonding Kinds of atoms involved?

4. Ionic Bond 1. Electrostatic force of attraction between oppositely charged ions 2. Ions result from the transfer of one or more electrons from a metal to a nonmetal (Trans of NaCl) 3. Why do metals lose electrons to form cations? 4. Why do nonmetals gain electrons to form anions?

5. Figure 9.1

6. Conditions Needed for Ionic Bond Formation 1. Chemical Bonding occurs only if it results in a decrease in PE » i.e. The process is exothermic 2. Cation formation is Endothermic (PE increases)....Why? » Relate to Ionization Energy 3. Anion formation is Exothermic (PE decreases)......Why? » Relate to Electron Affinity

7. Conditions Needed for Ionic Bond Formation 1. Cation formation is usually more endothermic than Anion formation is exothermic » Why then is Ionic Bond formation EXOTHERMIC?

8. Must Consider Lattice Energy 1. Lattice Energy » PE lowering due to the attraction of anions to cations » Highly Exothermic 2. Ionic bonding will only result when...... » Lattice Energy is more exothermic than E. A. + I.E. is endothermic E.g Li (s) + ½ F 2 (g)  LiF (s)

9. Figure 9.6

11. Figure 9.7

12. Factors that affect Lattice Energy 1. Lattice energy a.Depends on the charge, size and distance between the ions involved— Why?? b.Due to the electrostatic attractions between cations and anions  Electrostatic attractions depends on… Charge and size of ions—Why? Distance between ions—Why?

14. Periodic Trends in Lattice Energy 1. Down a group a.Down group IA b.Down group IIA c.Down group IIIA 2. Across a period a.Across period 2

16. Electron Configurations of Ions 1. Octet Rule Atoms of many elements tend to gain, lose, or share electrons until their valence shell contains 8 electrons

17. Rules for Writing Electron Configurations of Ions... 1. Group IA, IIA Metals and Aluminum »Lose electrons until reach Noble gas configuration 2. Nonmetals »Gain electrons until reach Noble gas configuration 3. Write the electron configurations for the ions in...... »KCl, CaCl 2, AlCl 3, CaO, Na 2 O, Al 2 O 3

18. Rules for Writing Electron Configurations of Ions... 1. Transition and Post-transition Metals » Do NOT obey the Octet Rule!! » More than one ion is often possible 2. Transition Metals » Lose s-Sublevel electrons, then d-electrons e.g. Fe 2+, Fe 3+, Zn 2+, Cu 1+, Cu 2+, 3. Post Transition Metals » Lose p-sublevel electrons, then s-electrons e.g. Sn 2+, Sn 4+, Pb 2+, Pb 4+

19. Lewis Symbols 1. Symbol of element surrounded by valence electrons »Used to represent bond formation 2. Write Lewis Symbols for.... »Representative Elements, Groups IA - VIIA Note: Group Number = number of valence electrons

21. Using Lewis Symbols to Illustrate Ionic Bond Formation 1. Use Lewis Symbols to diagram the reaction that produces the following compounds..... » KCl, CaCl 2, AlCl 3, CaO, Na 2 O, Al 2 O 3 » ZnCl 2

22. Explaining the Properties of Ionic compounds 1. Ionic compounds a.Have high melting points and boiling points (all are solids at room temp.) b.Hard, but brittle solids c.Conduct electricity in as liquids, but not as solids

27. Covalent Bonding 1. Involve the sharing of one or more PAIRS of electrons between atoms of nonmetallic elements 2. Occurs when ionic bond formation is not favored energetically »i.e. when.... I.E. + E.A. is more endothermic than the lattice energy is exothermic

28. Bond formation between two Hydrogen Atoms a)Large distance between atoms b.Atoms approach each other c.Covalent bond formation HHHHH2H2

29. Bond Length 1. Determined by a balance between the following...... a)Attractions of shared electrons to both nuclei –Causes a decrease in PE b)Repulsion between both nuclei –Causes an increase in PE

30. Figure 9.12

31. Figure 9.11

32. Figure 9.13

34. Bond Energy 1. Amount of energy released during bond formation 2. Amount of energy needed to break a bond

39. In Quartz: each Si atom is covalently bonded to 4 O atom. Each O atom is bonded to 2 Si atoms In Diamond: each C atom is covalently bonded to 4 other C atoms. Fig. 9.15 Network Covalent solids have very high melting points

40. Illustrating Covalent Bonding with Lewis Structures 1. Apply the Octet Rule »Atoms tend to share electrons until their valence shell contains 8 electrons 2. Use Lewis Structures to illustrate bond formation for..... »H 2, F 2, H 2 O, NH 3, CH 4 3. Multiple Bonds » N 2, SiO 2, NO 3 -

41. Guidelines for writing Lewis Structures 1. Decide which atoms are bonded 2. Count all valence electrons 3. Place 2 electrons in each bond 4. Complete the octets of the atoms attached to the central atom by adding electrons in pairs 5. Place any remaining electrons on the central atom in pairs 6. If the central atom does not have an octet, form double bonds, or if necessary, a triple bond.

43. Nonpolar vs Polar Covalent Bonding 1. Nonpolar Covalent Bond »Involves equal sharing of an electron pair between two nuclei –Pure nonpolar bonds are quite uncommon....Why?? 2. Polar Covalent Bond »Unequal sharing of electrons –Results from the electronegativity difference between atoms of different elements

44. Figure 9.16

45. Figure 9.17

47. Electronegativity Differences and Bond Types 1. Pure Nonpolar Covalent: 0 2. More Nonpolar than Polar: < 0.5 3. Polar Covalent: ~ 0.5 to 1.7 4. More Ionic than Polar Covalent: > 1.7

50. Some Examples 1. Indicate the kind of bonding in..... a)Water b)Ammonia c)Carbon dioxide d)Aluminum Chloride e)Methane f)Fatty Acids

51. Polar Bonds vs Polar Molecules 1. Why are water molecules polar, whereas carbon dioxide molecules are nonpolar?

52. Figure 9.21 Properties of the Period 3 chlorides.

53. Explaining the Properties of Metals a.Have high melting points (all but Hg are solids at room temp.) b.Malleable (deform when a force is applied) c.Conduct electricity

54. Figure 9.24 metal is deformed The reason metals deform. Why metals deform: Metal atoms slide past each other when a force is applied Why do metals conduct electricity? Figure 9.24 Explaining the Properties of Metals

55. Table 9.5 Melting and Boiling Points of Some Metals Elementmp( 0 C)bp( 0 C) Lithium (Li)1801347 Tin (Sn)2322623 Aluminum (Al)6602467 Barium (Ba)7271850 Silver (Ag)9612155 Copper (Cu)10832570 Uranium (U)11303930

56. Melting points of the Group 1A(1) and Group 2A(2) elements. Figure 9.23

57. Tools of the Laboratory: Infrared Spectroscopy Figure B9.1 Some vibrational modes in a diatomic molecule

58. Tools of the Laboratory: Infrared Spectroscopy Figure B9.1 Some vibrational modes in a triatomic molecule

59. Tools of the Laboratory: Infrared Spectroscopy Figure B9.1 The infrared (IR) spectrum of acrylonitrile.