1. Chapter Eight Comparing diamond & graphite: The bounding of substances (chemical) has a profound effect on chemical and physical properties. Comparing silicon & carbon: Group 4A CO 2 vs SiO 2 Molecular bonding and structure play the central role in understand the properties of matter and in the determining the course of all chemical reactions. Object: Bonding and Structure of materials (chemicals)

2. 8–28–2 Questions to Consider What is meant by the term “chemical bond?” Why do atoms bond with each other to form molecules? How do atoms bond with each other to form molecules?

3. 8–38–3 Figure 10.27: Examples of silicate anions, all of which are based on SiO 4 4- tetrahedra.

4. 8–48–4 Key Ideas in Bonding Ionic Bonding: Electrons are transferred, NaCl, Ionic compounds, Coulomb’s force Covalent Bonding: Electrons are shared equally, H 2 What about intermediate cases? –HF, polar covalent bonding Chemical Bond: energy, force ＃ : intramolecular and intermolecular bond Types of Chemical Bonds:

5. 8–58–5 Figure 8.1 a & b (a) The Interaction of Two Hydrogen Atoms (b) Energy Profile as a Function of the Distance Between the Nuclei of the Hydrogen Atoms

6. 8–68–6 Figure 8.2 The Effect of an Electric Field on Hydrogen Fluoride Molecules

7. 8–78–7 If lithium and fluorine react, which has more attraction for an electron? Why? In a bond between fluorine and iodine, which has more attraction for an electron? Why? 8.2 Electronegativity

8. 8–88–8 Figure 8.3 The Pauling Electronegativity Vaules Δ = (H–X) act - (H–X) exp

9. 8–98–9 Table 8.1 The Relationship Between Electronegativity and Bond Type

10. 8–10 Figure 8.4 An Electrostatic Potential Map of HF 8.3 Bond Polarity and Dipole Moments

11. 8–11 Figure 8.5 a-c The Charge Distribution in the Water Molecule

12. 8–12 Figure 8.6 a-c The Structure and Charge Distribution of the Ammonia Molecule

13. 8–13 Figure 8.7 a-c The Carbon Dioxide Molecule

14. 8–14 Table 8.2 Types of Molecules with Polar Bonds but No Resulting Dipole Moment

15. 8–15 e.p. Diagram HCL

16. 8–16 e.p.Diagram SO 3

17. 8–17 e.p. Diagram CH 4

18. 8–18 e.p. Diagram H 2 S

19. 8–19 Question Which of the following bonds would be the least polar yet still be considered polar covalent? Mg-O C-O O-O Si-O N-O

20. 8–20 Ions: Electron Configurations and Sizes What we can “read” from the periodic table: Trends for –Atomic size –Ion radius –Ionization energy –Electronegativity Electron configurations Predicting formulas for ionic compounds Ranking polarity of covalent bonds

21. 8–21 Table 8.3 Common Ions with Noble Gas Configurations in Ionic Compounds

22. 8–22 Figure 8.8 Sizes of Ions Related to Positions of the Elements on the Periodic Table 8.4 Ions: Electron Configurations and Sizes

23. 8–23 Arrange the following, without consulting any specific listed radii, in order from smallest to largest size (radius) S 2– ; Cl – ; K + ; Rb + 1.S 2– ; Cl – ; K + ; Rb + 2.K + ; Rb + ; Cl – ; S 2– 3.S 2– ; Cl – ; Rb +; ; K + 4.S 2– ; Cl – ; Rb + ; K + QUESTION

24. 8–24 Formation of an Ionic Solid 1. Sublimation of the solid metal M(s)  M(g) [endothermic] 2.Ionization of the metal atoms M(g)  M + (g) + e  [endothermic] 3.Dissociation of the nonmetal 1 / 2 X 2 (g)  X(g) [endothermic] 4. Formation of X  ions in the gas phase: X(g) + e   X  (g) [exothermic] 5. Formation of the solid MX M + (g) + X  (g)  MX(s) [quite exothermic]

25. 8–25 Figure 8.9 The Energy Changes Involved in the Formation of Lithium Fluoride from Its Elements

26. 8–26 Figure 8.11 Comparison of the Energy Changes Involved in the Formation of Solid Sodium Fluoride and Solid Magnesium Oxide

27. 8–27 Figure 8.10 a & b The Structure of Lithium Fluoride

28. 8–28 Figure 8.12 a-c The Three Possible Types of Bonds

29. 8–29 Figure 8.13 The Relationship Between the Ionic Character of a Covalent Bond and the Electronegativity Difference of the Bounded Atoms

30. 8–30 Models Models are attempts to explain how nature operates on the microscopic level based on experiences in the macroscopic world.

31. 8–31 Fundamental Properties of Models 1.A model does not equal reality. 2.Models are oversimplifications, and are therefore often wrong. 3.Models become more complicated as they age. 4.We must understand the underlying assumptions in a model so that we don’t misuse it.

32. 8–32 Model of Chemical Bond Individual bond occurs between pair electron. Average individual bond energy Bond energy values can used to calculate the reaction energies.

33. 8–33 Table 8.4 Average Bond Energies (kj/mol)

34. 8–34 Table 8.5 Bond Lengths for Selected Bonds

35. 8–35 Localized Electron Model A molecule is composed of atoms that are bound together by sharing pairs of electrons using the atomic orbitals of the bound atoms. bonding pairs, lone pairs

36. 8–36 Localized Electron Model 1.Description of valence electron arrangement (Lewis structure). 2.Prediction of geometry (VSEPR model). 3.Description of atomic orbital types used to share electrons or hold long pairs.

37. 8–37 Lewis Structure Shows how valence electrons are arranged among atoms in a molecule. Reflects central idea that stability of a compound relates to noble gas electron configuration.

38. 8–38 Lewis Structures 1.Sum the valence electrons. 2.Place bonding electrons between pairs of atoms. 3.Atoms usually have noble gas configurations.

39. 8–39 5)

40. 8–40 Resonance More than one Lewis structure for a given molecule. Ex. NO 3 - Resonance structure, Ex. NO 2 - Formal charge: (# of VEs on free atom)-(# of VEs assigned to the atom in the molecule), Ex. SO 4 2- Rules of formal charge Assigning the structure based on formal charge (predicted structure should be judged by expt.

41. 8–41

42. 8–42 Molecular Structure: VSEPR Model Molecular structure: 3D arrangement of the atoms in a molecule. VSEPR: valence shell electron-pair repulsion Assume: The structure around a given atom is determined principally by minimizing electron pair repulsions. (Bonding and nonbonding pairs around a given atom should be positioned as far as possible.) Ex. BeCl 2, BF 3, CH 4

43. 8–43 Figure 8.15 The Molecular Structure of Methane

44. 8–44 Balloons Tied Together Naturally Form Tetrahedral Shape

45. 8–45 Predicting a VSEPR Structure 1.Draw Lewis structure. 2.Put pairs as far apart as possible. 3.Determine positions of atoms from the way electron pairs are shared. 4.Determine the name of molecular structure from positions of the atoms.

46. 8–46 Figure 8.16 a-c The Molecular Structure of Ammonia is a Trigonal Pyramid

47. 8–47 Figure 8.17 a-c The Tetrahedral Arrangement of Oxygen In a Water Molecule

48. 8–48 Figure 8.18 The Bond Angles In the CH 4, NH 3, and H 2 O Molecules

49. 8–49 Figure 8.19 a & b In a Bonding Pair of Electrons the Electrons are Shared by Two Nuclei (b) In a Lone Pair, Both Electrons Must Be Close to a Single Nucleus

50. 8–50 Predicting a VSEPR Structure 1.Draw Lewis structure. 2.Put pairs as far apart as possible. 3.Determine positions of atoms from the way electron pairs are shared. 4.Determine the name of molecular structure from positions of the atoms. 5.Lone pairs require more room than bonding pairs

51. 8–51 Table 8.6 Arrangements of Electron Pairs Around an Atom Yielding Minimum Repulsion

52. 8–52 Table 8.7 Structures of Molecules that Have Four Electron Pairs Around the Central Atom

53. 8–53 Table 8.8 Structures of Molecules with Five Electron Pairs Around the Central Atom

54. 8–54 Determine the shape for each of the following molecules, and include bond angles: HCN PH 3 SF 4 O 3 KrF 4

55. 8–55 Figure 8.20 a & b Possible Electron Pair Arrangements for XeF 4

56. 8–56 Figure 8.21 a-c Three Possible Arrangements of the Electron Pairs in the I 3 - Ion

57. 8–57 Figure 8.22 a-c The Molecular Structure of Methanol

58. 8–58 Queen Bee

59. 8–59 NH3

60. 8–60 PH3