1. 1 Quartz grows in beautiful, regular crystals.

2. 2 Lithium fluoride

3. 3 Bonds Forces that hold groups of atoms together and make them function as a unit.

4. 4 Bond Energy 4 It is the energy required to break a bond. 4 It gives us information about the strength of a bonding interaction.

5. 5 Bond Length The distance where the system energy is a minimum.

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. 7 Ionic Bonds Q 1 and Q 2 = numerical ion charges r = distance between ion centers (in nm)

8. 8 (a) The interaction of two hydrogen atoms. (b) Energy profile as a function of the distance between the nuclei of the hydrogen atoms.

9. 9 Electronegativity The ability of an atom in a molecule to attract shared electrons to itself.  = (H  X) actual  (H  X) expected

10. 10

11. 11 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.

12. 12 The effect of an electric field on hydrogen fluoride molecules. (a) When no electric field is present, the molecules are randomly oriented. b) When the field is turned on, the molecules tend to line up with their negative ends toward the positive pole and their positive ends toward the negative pole.

13. 13 (a) The charge distribution in the water molecule. (b) The water molecule in an electric field.

14. 14 (a) The structure and charge distribution of the ammonia molecule. The polarity of the N—H bonds occurs because nitrogen has a greater electronegativity than hydrogen. (b) The dipole moment of the ammonia molecule oriented in an electric field.

15. 15 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.

16. 16 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

17. 17 (a) The carbon dioxide molecule. (b) The opposed bond polarities cancel out, and the carbon dioxide has no dipole moment.

18. 18

19. 19

20. 20 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.

21. 21 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]

22. 22 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]

23. 23 The energy changes involved in the formation of solid lithium fluoride from its elements.

24. 24 The structure of lithium fluoride.

25. 25 Comparison of the energy changes involved in forming solid sodium fluoride and solid magnesium oxide.

26. 26

27. 27 The Pauling electronegativity values. Electronegativity generally increases across a period and decreases down a group.

28. 28 The three possible types of bonds (a) a covalent bond formed between identical atoms (b) a polar covalent bond, with both ionic and covalent components; and (c) an ionic bond with no electron sharing.

29. 29 The relationship between the ionic character of a covalent bond and the electronegativity difference of the bonded atoms.

30. 30 Molten NaCl conducts an electric current, indicating the presence of moblie Na+ and Cl- ions.

31. 31 Sizes of ions related to positions of the elements on the periodic table.

32. 32 Q 1, Q 2 = charges on the ions r = shortest distance between centers of the cations and anions

33. 33 Models Models are attempts to explain how nature operates on the microscopic level based on experiences in the macroscopic world.

34. 34 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. 4 We must understand the underlying assumptions in a model so that we don’t misuse it.

35. 35 Bond Energies Bond breaking requires energy (endothermic). Bond formation releases energy (exothermic).  H =  D( bonds broken )   D( bonds formed ) energy requiredenergy released

36. 36

37. 37 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.

38. 38 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.

39. 39 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.

40. 40 The molecular structure of methane. The tetrahedral arrangement of electron pairs produces a tetrahedral arrangement of hydrogen atoms.

41. 41 (a) The tetrahedral arrangement of electron pairs around the nitrogen atom in the ammonia molecule. (b) Three of the electron pairs around nitrogen are shared with hydrogen atoms as shown and one is a lone pair. Although the arrangement of electron pairs is tetrahedral, as in the methane molecule, the hydrogen atoms in the ammonia molecule occupy only three corners of the tetrahedron. A lone pair occupies the fourth corner. (c) Note that molecular geometry is trigonal pyramidal, not tetrahedral.

42. 42 (a) The tetrahedral arrangement of the four electron pairs around oxygen in the water molecule. (b) Two of the electron pairs are shared between oxygen and the hydrogen atoms and two are lone pairs. (c) The V-shaped molecular structure of the water molecule.

43. 43 The bond angles in the CH 4, NH 3, and H 2 O molecules.

44. 44 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.

45. 45 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.

46. 46 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

47. 47 Formal Charge Not as good Better

48. 48 VSEPR Model The structure around a given atom is determined principally by minimizing electron pair repulsions.

49. 49 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.

50. 50 Figure 8.18: (a) 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 and tend to take up more of the space around that atom.

51. 51

52. 52 Molecular structure of PCl 6 -

53. 53 Octahedral electron arrangement for Xe

54. 54 Possible electron-pair arrangements for XeF 4.

55. 55 Three possible arrangements of the electron pairs in the I 3 - ion.

56. 56 The molecular structure of methanol. (a) The arrangement of electron pairs and atoms around the carbon atom. (b) The arrangement of bonding and lone pairs around the oxygen atom. (c) The molecular structure.