1. Copyright©2000 by Houghton Mifflin Company. All rights reserved. 1 Bonds Forces that hold groups of atoms together and make them function as a unit.

2. Copyright©2000 by Houghton Mifflin Company. All rights reserved. 2 Bond Energy 4 It is the energy required to break a bond. 4 It gives us information about the strength of a bonding interaction.

3. Copyright©2000 by Houghton Mifflin Company. All rights reserved. 3 Bond Length The distance where the system energy is a minimum.

4. Copyright©2000 by Houghton Mifflin Company. All rights reserved. 4 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.

5. Copyright©2000 by Houghton Mifflin Company. All rights reserved. 5 Ionic Bonds Q 1 and Q 2 = numerical ion charges r = distance between ion centers (in nm)

6. Copyright©2000 by Houghton Mifflin Company. All rights reserved. 6 Electronegativity The ability of an atom in a molecule to attract shared electrons to itself.  = (H  X) actual  (H  X) expected

7. Copyright©2000 by Houghton Mifflin Company. All rights reserved. 7 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.

8. Copyright©2000 by Houghton Mifflin Company. All rights reserved. 8 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.

9. Copyright©2000 by Houghton Mifflin Company. All rights reserved. 9 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

10. Copyright©2000 by Houghton Mifflin Company. All rights reserved. 10 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.

11. Copyright©2000 by Houghton Mifflin Company. All rights reserved. 11 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]

12. Copyright©2000 by Houghton Mifflin Company. All rights reserved. 12 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]

13. Copyright©2000 by Houghton Mifflin Company. All rights reserved. 13 Q 1, Q 2 = charges on the ions r = shortest distance between centers of the cations and anions

14. Copyright©2000 by Houghton Mifflin Company. All rights reserved. 14 Models Models are attempts to explain how nature operates on the microscopic level based on experiences in the macroscopic world.

15. Copyright©2000 by Houghton Mifflin Company. All rights reserved. 15 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.

16. Copyright©2000 by Houghton Mifflin Company. All rights reserved. 16 Bond Energies Bond breaking requires energy (endothermic). Bond formation releases energy (exothermic).  H =  D( bonds broken )   D( bonds formed ) energy requiredenergy released

17. Copyright©2000 by Houghton Mifflin Company. All rights reserved. 17 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.

18. Copyright©2000 by Houghton Mifflin Company. All rights reserved. 18 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.

19. Copyright©2000 by Houghton Mifflin Company. All rights reserved. 19 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.

20. Copyright©2000 by Houghton Mifflin Company. All rights reserved. 20 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.

21. 21 Writing Lewis Formulas: The Octet Rule N - A = S rule Simple mathematical relationship to help us write Lewis dot formulas. N = number of electrons needed to achieve a noble gas configuration. N usually has a value of 8 for representative elements. N has a value of 2 for H atoms. A = number of electrons available in valence shells of the atoms. A is equal to the periodic group number for each element. A is equal to 8 for the noble gases. S = number of electrons shared in bonds. A-S = number of electrons in unshared, lone, pairs.

22. 22 Writing Lewis Formulas: The Octet Rule For ions we must adjust the number of electrons available, A. Add one e - to A for each negative charge. Subtract one e - from A for each positive charge. The central atom in a molecule or polyatomic ion is determined by: The atom that requires the largest number of electrons to complete its octet goes in the center. For two atoms in the same periodic group, the less electronegative element goes in the center.

23. 23 Writing Lewis Formulas: The Octet Rule Example 7-2:Write Lewis dot and dash formulas for hydrogen cyanide, HCN. N = 2 (H) + 8 (C) + 8 (N) = 18 A = 1 (H) + 4 (C) + 5 (N) = 10 S = 8 A-S = 2 This molecule has 8 electrons in shared pairs and 2 electrons in lone pairs.

24. 24 Writing Lewis Formulas: The Octet Rule Example 7-3:Write Lewis dot and dash formulas for the sulfite ion, SO 3 2-. N = 8 (S) + 3 x 8 (O) = 32 A = 6 (S) + 3 x 6 (O) + 2 (- charge) = 26 S = 6 A-S = 20 Thus this polyatomic ion has 6 electrons in shared pairs and 20 electrons in lone pairs. Which atom is the central atom in this ion? You do it!

25. 25 Writing Lewis Formulas: The Octet Rule What kind of covalent bonds, single, double, or triple, must this ion have so that the six shared electrons are used to attach the three O atoms to the S atom?

26. LEWIS STRUCTURES 1 : draw skeleton of species 2 : count e - in species 3 : subtract 2 e - for each bond in skeleton 4 : distribute remaining e -

27. Copyright©2000 by Houghton Mifflin Company. All rights reserved. 27 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.

28. 28 Resonance Example 7-4:Write Lewis dot and dash formulas for sulfur trioxide, SO 3. You do it! You do it! N = 8 (S) + 3 x 8 (O) = 32 A = 6 (S) + 3 x 6 (O) = 24 S = 8 A-S= 16

29. 29 Resonance There are three possible structures for SO 3. The double bond can be placed in one of three places. o oWhen two or more Lewis formulas are necessary to show the bonding in a molecule, we must use equivalent resonance structures to show the molecule’s structure. o oDouble-headed arrows are used to indicate resonance formulas.

30. 30 Resonance Resonance is a flawed method of representing molecules. There are no single or double bonds in SO 3. In fact, all of the bonds in SO 3 are equivalent. The best Lewis formula of SO 3 that can be drawn is:

31. RESONANCE u a condition occurring when more than one valid Lewis structure can be written for a particular species

32. RESONANCE The actual electronic structure is not represented by any one of the Lewis structures; it is an average of them

33. 33 Writing Lewis Formulas: Limitations of the Octet Rule There are some molecules that violate the octet rule. For these molecules the N - A = S rule does not apply: 1. The covalent compounds of Be. 2. The covalent compounds of the IIIA Group. 3. Species which contain an odd number of electrons. 4. Species in which the central element must have a share of more than 8 valence electrons to accommodate all of the substituents. 5. Compounds of the d- and f-transition metals.

34. 34 Writing Lewis Formulas: Limitations of the Octet Rule In those cases where the octet rule does not apply, the substituents attached to the central atom nearly always attain noble gas configurations. The central atom does not have a noble gas configuration but may have fewer than 8 (exceptions 1, 2, & 3) or more than 8 (exceptions 4 & 5).

35. 35 Writing Lewis Formulas: Limitations of the Octet Rule Example 7-5: Write dot and dash formulas for BBr 3. This is an example of exception #2. You do it!

36. 36 Writing Lewis Formulas: Limitations of the Octet Rule Example 7-6: Write dot and dash formulas for AsF 5. You do it!

37. Copyright©2000 by Houghton Mifflin Company. All rights reserved. 37 VSEPR Model The structure around a given atom is determined principally by minimizing electron pair repulsions.

38. Copyright©2000 by Houghton Mifflin Company. All rights reserved. 38 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.

39. Distinguish between ELECTRONIC Geometry & MOLECULAR Geometry

40. CH 4 Electronic geometry: tetrahedral Molecular geometry: tetrahedral Bond angle = 109.5 0

41. H3O+H3O+ Electronic geometry: tetrahedral Molecular geometry: trigonal pyramidal Bond angle ~ 107 0

42. H2OH2O Electronic geometry: tetrahedral Molecular geometry: bent Bond angle ~ 104.5 0

43. NH 2 - Electronic geometry: tetrahedral Molecular geometry: bent Bond angle ~ 104.5 0

44. “Octet Rule” holds for connecting atoms, but may not for the central atom.

45. BF 3 Electronic geometry: trigonal planar Molecular geometry: trigonal planar Bond angle =120 0

46. PF 5 Electronic geometry: trigonal bipyramidal Molecular geometry: trigonal bipyramidal Bond angle = 120 0 & 90 0

47. SF 4 Electronic geometry: trigonal bipyramidal Molecular geometry: see-saw Bond angle = 120 0 & 90 0

48. ICl 3 Electronic geometry: trigonal bipyramidal Molecular geometry: T-shape Bond angle <= 90 0

49. I3-I3- Electronic geometry: trigonal bipyramid Molecular geometry: linear Bond angle = 180 0

50. PCl 6 - Electronic geometry: octahedral Molecular geometry: octahedral Bond angle = 90 0

51. BrF 5 Electronic geometry: octahedral Molecular geometry: square pyramidal Bond angle ~ 90 0

52. ICl 4 - Electronic geometry: octahedral Molecular geometry: square planar Bond angle = 90 0

53. What happens when there is not enough electrons to “satisfy” the central atom? Double or Triple Bonds are needed.

54. EXAMPLES  Ethene  Acetic Acid  Oxygen  Nitrogen

55. Copyright©2000 by Houghton Mifflin Company. All rights reserved. 55 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

56. FORMAL CHARGE u the charge assigned to an atom in a molecule or polyatomic ion

57. Formal charge = #VE on free atom - #VE assigned to atom in Lewis Structure

58. Assignment of e - 1 1. Lone pairs belong entirely to atom in question 2 2. Shared e - are divided equally between the two sharing atoms

59. Copyright©2000 by Houghton Mifflin Company. All rights reserved. 59 Formal Charge Not as good Better

60. EXAMPLES u Nitrate ion u Ozone

61. The sum of the formal charges of all atoms in a species must equal the overall charge on the species.

62. If nonequivalent Lewis structures exist, the one(s) that best describe(s) the bonding in the species has... FAVORED LEWIS STRUCTURES

63. 1 1. formal charges closest to zero 2 2. negative formal charge is on the most electronegative atom

64. EXAMPLES u Carbon dioxide u Thiocyanate ion u Sulfate ion