Chapter 9 (Silberberg 3 rd Edition) Models of Chemical Bonding

1. Chapter 9 (Silberberg 3rd 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 • What’s a Chemical Bond? • Attraction that holds atoms or ions together in compounds • Ionic Bonding vs Covalent Bonding • What’s the difference? • Kinds of atoms involved? • Metallic Bonding • Kinds of atoms involved?

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

5. Figure 9.1

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

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

8. Must Consider Lattice Energy • Lattice Energy • PE lowering due to the attraction of anions to cations • Highly Exothermic • Ionic bonding will only result when...... • Lattice Energy is more exothermic than E. A. + I.E. is endothermic E.g Li (s) + ½ F2 (g)  LiF (s)

9. Figure 9.6

11. Figure 9.7

12. Factors that affect Lattice Energy • Lattice energy • Depends on the charge, size and distance between the ions involved— Why?? • 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 • Down a group • Down group IA • Down group IIA • Downgroup IIIA • Across a period • Across period 2

16. Electron Configurations of Ions • 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... • Group IA , IIA Metals and Aluminum • Lose electrons until reach Noble gas configuration • Nonmetals • Gain electrons until reach Noble gas configuration • Write the electron configurations for the ions in...... • KCl, CaCl2, AlCl3, CaO, Na2 O, Al2O3

18. Rules for Writing Electron Configurations of Ions... • Transition and Post-transition Metals • Do NOT obey the Octet Rule!! • More than one ion is often possible • Transition Metals • Lose s-Sublevel electrons, then d-electrons e.g. Fe 2+, Fe 3+ , Zn 2+ , Cu1+ , Cu2+ , • Post Transition Metals • Lose p-sublevel electrons, then s-electrons e.g. Sn 2+ , Sn 4+ , Pb 2+ , Pb 4+

19. Lewis Symbols • Symbol of element surrounded by valence electrons • Used to represent bond formation • 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 • Use Lewis Symbols to diagram the reaction that produces the following compounds..... • KCl, CaCl2, AlCl3, CaO, Na2 O, Al2O3 • ZnCl2

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

27. Covalent Bonding • Involve the sharing of one or more PAIRS of electrons between atoms of nonmetallic elements • 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 H H H H H2 • Atoms approach each other • Covalent bond formation • Large distance between atoms

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

30. Figure 9.12

31. Figure 9.11

32. Figure 9.13

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

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

40. Illustrating Covalent Bonding with Lewis Structures • Apply the Octet Rule • Atoms tend to share electrons until their valence shell contains 8 electrons • Use Lewis Structures to illustrate bond formation for..... • H2, F2, H2O, NH3, CH4 • Multiple Bonds • N2, SiO2 , NO3-

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

43. Nonpolar vs Polar Covalent Bonding • Nonpolar Covalent Bond • Involves equal sharing of an electron pair between two nuclei • Pure nonpolar bonds are quite uncommon....Why?? • 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 • Pure Nonpolar Covalent: 0 • More Nonpolar than Polar: < 0.5 • Polar Covalent: ~ 0.5 to 1.7 • More Ionic than Polar Covalent: > 1.7

50. Some Examples • Indicate the kind of bonding in..... • Water • Ammonia • Carbon dioxide • Aluminum Chloride • Methane • Fatty Acids