1. Copyright©2000 by Houghton Mifflin Company. All rights reserved. 1 Chemistry FIFTH EDITION by Steven S. Zumdahl University of Illinois

2. Copyright©2000 by Houghton Mifflin Company. All rights reserved. 2 Chemistry FIFTH EDITION Chapter 8 Chemical Foundations Dr. JGS RAMON

3. Copyright©2000 by Houghton Mifflin Company. All rights reserved. 3

4. 4 Bonds Forces that hold groups of atoms together and make them function as a unit.

5. Forces of Attraction and Repulsion Copyright©2000 by Houghton Mifflin Company. All rights reserved. 5 attraction attraction attraction attraction repulsions

6. 6 Ionic Bonds 4 Formed from electrostatic attractions of closely packed, oppositely charged ions. 4 Formed when an atom that easily loses electrons reacts with one that has a high electron affinity.

7. Copyright©2000 by Houghton Mifflin Company. All rights reserved. 7 Ionic Bonds Q 1 and Q 2 = numerical ion charges r = distance between ion centers (in nm)

8. Copyright©2000 by Houghton Mifflin Company. All rights reserved. 8 Figure 8.1 Interaction of Two H Atoms and the Energy Profile

9. Copyright©2000 by Houghton Mifflin Company. All rights reserved. 9 Types of Bonds Types of Bond CovalentIonic PolarNon-Polar

10. Copyright©2000 by Houghton Mifflin Company. All rights reserved. 10 Technically all bonds are covalent!!! Its just a matter of degree of covalency that changes

11. Copyright©2000 by Houghton Mifflin Company. All rights reserved. 11 Bond Energy 4 It is the energy required to break a bond. 4 It gives us information about the strength of a bonding interaction.

12. Copyright©2000 by Houghton Mifflin Company. All rights reserved. 12 Bond Length The distance where the system energy is a minimum.

13. Copyright©2000 by Houghton Mifflin Company. All rights reserved. 13 Electronegativity The ability of an atom in a molecule to attract shared electrons to itself.  = (H  X) actual  (H  X) expected

14. Copyright©2000 by Houghton Mifflin Company. All rights reserved. 14 Figure 8.3 The Pauling Electronegativity Values

15. Copyright©2000 by Houghton Mifflin Company. All rights reserved. 15 Figure 8.11 Three Possible Types of Bonds

16. Copyright©2000 by Houghton Mifflin Company. All rights reserved. 16 Figure 8.12 The Relationship Between the Ionic Character of a Covalent Bond and the Electronegativity Difference of the Bonded Atoms

17. Electronegativity Difference Ionic Character (%) Covalent Character (%) Bond Type 0.00100Covalent 0.5595Covalent 1.02080Covalent 1.54060Polar 2.06040Polar 2.57525Ionic 3.09010Ionic

18. Copyright©2000 by Houghton Mifflin Company. All rights reserved. 18 Polarity A molecule, such as HF, that has a center of positive charge and a center of negative charge is said to be polar, or to have a dipole moment.

19. Copyright©2000 by Houghton Mifflin Company. All rights reserved. 19 Figure 8.2 The Effect of an Electric Field on Hydrogen Fluoride Molecules

20. Relative Bond Polarity Copyright©2000 by Houghton Mifflin Company. All rights reserved. 20 Sample Exercise 8.1, P.355 Order the following bonds according to increasing polarity: H-H, O-H, Cl-H, S-H, and F-H H-H < S-H < Cl-H < O-H < F-H Exercise 25, P.406 Exercise 26, P.406

21. Copyright©2000 by Houghton Mifflin Company. All rights reserved. 21 Figure 8.4 Dipole Moment for H 2 O

22. Copyright©2000 by Houghton Mifflin Company. All rights reserved. 22 Figure 8.5 Dipole Moment for NH 3

23. Copyright©2000 by Houghton Mifflin Company. All rights reserved. 23 Figure 8.6 (a) Carbon Dioxide (b) Opposed Bond Polarities

24. Relative Bond Polarity Copyright©2000 by Houghton Mifflin Company. All rights reserved. 24 Sample Exercise 8.1, P.355 Order the following bonds according to increasing polarity: H-H, O-H, Cl-H, S-H, and F-H H-H < S-H < Cl-H < O-H < F-H Exercise 25, P.406 Exercise 26, P.406

25. Copyright©2000 by Houghton Mifflin Company. All rights reserved. 25 Achieving Noble Gas Electron Configurations (NGEC) Two nonmetals react: They share electrons to achieve NGEC. A nonmetal and a representative group metal react (ionic compound): The valence orbitals of the metal are emptied to achieve NGEC. The valence electron configuration of the nonmetal achieves NGEC.

26. Copyright©2000 by Houghton Mifflin Company. All rights reserved. 26 Isoelectronic Ions Ions containing the the same number of electrons (O 2 , F , Na +, Mg 2+, Al 3+ ) O 2  > F  > Na + > Mg 2+ > Al 3+ largest smallest

27. Copyright©2000 by Houghton Mifflin Company. All rights reserved. 27 Figure 8.7 Sizes of Ions Related to Positions of the Elements in the Periodic Table

28. Copyright©2000 by Houghton Mifflin Company. All rights reserved. 28 Models Models are attempts to explain how nature operates on the microscopic level based on experiences in the macroscopic world.

29. Copyright©2000 by Houghton Mifflin Company. All rights reserved. 29 Fundamental Properties of Models 4 A model does not equal reality. 4 Models are oversimplifications, and are therefore often wrong. 4 Models become more complicated as they age.  We must understand the underlying assumptions in a model so that we don ’ t misuse it.

30. Copyright©2000 by Houghton Mifflin Company. All rights reserved. 30 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.

31. Copyright©2000 by Houghton Mifflin Company. All rights reserved. 31 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.

32. Copyright©2000 by Houghton Mifflin Company. All rights reserved. 32 Lewis Structure 4 Shows how valence electrons are arranged among atoms in a molecule. 4 Reflects central idea that stability of a compound relates to noble gas electron configuration.

33. Copyright©2000 by Houghton Mifflin Company. All rights reserved. 33 Comments About the Octet Rule 4 2nd row elements C, N, O, F observe the octet rule. 4 2nd row elements B and Be often have fewer than 8 electrons around themselves - they are very reactive. 4 3rd row and heavier elements CAN exceed the octet rule using empty valence d orbitals. 4 When writing Lewis structures, satisfy octets first, then place electrons around elements having available d orbitals.

34. Copyright©2000 by Houghton Mifflin Company. All rights reserved. 34 Steps in Drawing Lewis Structure 1. Sum the valence electrons from all the atoms. NOTE: Don’t worry about keeping track of which electrons come from which atom. Just count the total valence electrons. 2. Identify the central atom. NOTE: Usually, but not always, this will be the least electronegative atom(s). 3. Use a pair of electrons to form bonds (called bonding pairs) between the central atom and the other atoms. 4. Complete the octet of the surrounding atoms by adding pairs of electrons (called lone pairs). 5. Complete the octet of the central atom by adding remaining electrons or by converting lone pairs to bonding pairs. NOTE the comments on the octet rule.

35. Drawing Lewis Structures Sample Exercise 8.6, P.379 Give the Lewis structure for each of the following. Copyright©2000 by Houghton Mifflin Company. All rights reserved. 35 a.HF b.N 2 c.NH 3 d.CH 4 e.CF 4 f.NO + Exercise 8.63 Exercise 8.64

36. Lewis Structures with Exceptions to the Octet Rule Sample Exercise 8.7, P.382 Write the Lewis structure for PCl 5. Copyright©2000 by Houghton Mifflin Company. All rights reserved. 36 Exercise 8.65 Exercise 8.66

37. Lewis Structures with Exceptions to the Octet Rule Sample Exercise 8.8, P.383 Write the Lewis structure for each molecule or ion. Copyright©2000 by Houghton Mifflin Company. All rights reserved. 37 Exercise 8.65 Exercise 8.66 a.ClF 3 b.XeO 3 c.RnCl 2 d.BeCl 2 e.ICl 4 -

38. Lewis Structures: One More Time Sample Exercise 8.9, P.385 Describe the electron arrangement in the nitrite anion using the localized electron model. Copyright©2000 by Houghton Mifflin Company. All rights reserved. 38 Exercise 8.67 Exercise 8.72

39. Copyright©2000 by Houghton Mifflin Company. All rights reserved. 39 Resonance Occurs when more than one valid Lewis structure can be written for a particular molecule. These are resonance structures. The actual structure is an average of the resonance structures.

40. Lewis Structures: One Last Thing Describe the electron arrangement in the sulfate anion using the localized electron model. Copyright©2000 by Houghton Mifflin Company. All rights reserved. 40

41. Copyright©2000 by Houghton Mifflin Company. All rights reserved. 41 Formal Charge The difference between the number of valence electrons (VE) on the free atom and the number assigned to the atom in the molecule. We need: 1.# VE on free neutral atom 2.# VE “ belonging ” to the atom in the molecule

42. Copyright©2000 by Houghton Mifflin Company. All rights reserved. 42 Formal Charge Not as good Better Formal Charge = valence e - - lone pair e - - ½ bonding pair e - For Carbon Dioxide: There are 2 possible Lewis structures. NOTE: Lower over-all formal charges leads to a more stable/plausible structure.

43. Lewis Structures with Formal Charges Sample Exercise 8.10, P.382 Give possible Lewis Structures for XeO 3, an explosive compound of Xenon. Which Lewis structure or structures are most appropriate according to the formal charges? Copyright©2000 by Houghton Mifflin Company. All rights reserved. 43 Exercise 8.75 Exercise 8.78

44. Copyright©2000 by Houghton Mifflin Company. All rights reserved. 44 VSEPR Model The structure around a given atom is determined principally by minimizing electron pair repulsions.

45. Copyright©2000 by Houghton Mifflin Company. All rights reserved. 45 VSEPR Model V alence S hell E lectron P air R epulsion Model (Still based on the localized electron model) Types of Electron Pairs: Bonding Pair (BP) – more localized Lone Pair (LP) – less localized BP-BP Repulsion < BP-LP Repulsion < LP-LP Repulsion

46. Copyright©2000 by Houghton Mifflin Company. All rights reserved. 46 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.

47. Molecular Structure Sample Exercise 8.11, P.394 Describe the molecular structure of the water molecule. Copyright©2000 by Houghton Mifflin Company. All rights reserved. 47 Exercise 8.83 Exercise 8.84

48. Copyright©2000 by Houghton Mifflin Company. All rights reserved. 48 Figure 8.17 The Bond Angles in the CH 4, NH 3, and H 2 O Molecules

49. Copyright©2000 by Houghton Mifflin Company. All rights reserved. 49 Figure 8.14 The Molecular Structure of Methane

50. Copyright©2000 by Houghton Mifflin Company. All rights reserved. 50 Figure 8.15 The Molecular Structure of NH 3

51. Copyright©2000 by Houghton Mifflin Company. All rights reserved. 51 Figure 8.16 The Molecular Structure of H 2 O

52. Copyright©2000 by Houghton Mifflin Company. All rights reserved. 52 Figure 8.17 The Bond Angles in the CH 4, NH 3, and H 2 O Molecules

53. Copyright©2000 by Houghton Mifflin Company. All rights reserved. 53 VSEPR Predicted Geometries

54. Copyright©2000 by Houghton Mifflin Company. All rights reserved. 54 Figure 8.20 Three Possible Arrangements of the Electron Pairs in the I 3 - Ion

55. Copyright©2000 by Houghton Mifflin Company. All rights reserved. 55 Figure 8.19 Possible Electron Pair Arrangements for XeF 4

56. Copyright©2000 by Houghton Mifflin Company. All rights reserved. 56 VSEPR Predicted Geometries

57. Copyright©2000 by Houghton Mifflin Company. All rights reserved. 57 Figure 8.21 The Molecular Structure of Methanol

58. Molecular Structure Sample Exercise 8.12, P.396 Copyright©2000 by Houghton Mifflin Company. All rights reserved. 58 Exercise 8.81 Exercise 8.82 Exercise 8.85 Exercise 8.86

59. Molecular Structure Sample Exercise 8.13, P.397 Copyright©2000 by Houghton Mifflin Company. All rights reserved. 59 Exercise 8.87 Exercise 8.90

60. Molecular Structure Sample Exercise 8.14, P.400 Copyright©2000 by Houghton Mifflin Company. All rights reserved. 60 Exercise 8.91 Exercise 8.92

61. Copyright©2000 by Houghton Mifflin Company. All rights reserved. 61 Why, Oh, why? Let’s look at methane, again:

62. Copyright©2000 by Houghton Mifflin Company. All rights reserved. 62 Bond Energies Bond breaking requires energy (endothermic). Bond formation releases energy (exothermic).  H =  D( bonds broken )   D( bonds formed ) energy requiredenergy released

63. Copyright©2000 by Houghton Mifflin Company. All rights reserved. 63 Figure 8.9 The Structure of Lithium Fluoride

64. Copyright©2000 by Houghton Mifflin Company. All rights reserved. 64 Lattice Energy The change in energy when separated gaseous ions are packed together to form an ionic solid. M + (g) + X  (g)  MX(s) Lattice energy is negative (exothermic) from the point of view of the system.

65. Copyright©2000 by Houghton Mifflin Company. All rights reserved. 65 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]

66. Copyright©2000 by Houghton Mifflin Company. All rights reserved. 66 Formation of an Ionic Solid (continued) 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]

67. Copyright©2000 by Houghton Mifflin Company. All rights reserved. 67 Q 1, Q 2 = charges on the ions r = shortest distance between centers of the cations and anions

68. Copyright©2000 by Houghton Mifflin Company. All rights reserved. 68 Figure 8.8 The Energy Changes Involved in the Formation of Solid Lithium Fluoride from its Elements

69. Copyright©2000 by Houghton Mifflin Company. All rights reserved. 69 Figure 8.10 Comparison of the Energy Changes Involved in the Formation of Solid Sodium Fluoride and Solid Magnesium Oxide