1. Bonding
Chapters 8 & 9

2. Bonds Ionic Covalent Metallic
Electrostatic attraction between oppositely charged ions
Covalent
Shared valence electrons
Metallic
Pooled electrons around a positive core

3. Octet Rule
Atoms will gain, lose, or share electrons to achieve a full valence shell
First energy level is limited to 2 electrons
d & f sublevels allow for more than 8

4. Ionic Bonding Electrostatic attraction due to oppositely charged ions
Lattice Energy
Energy released when gaseous ions form 1 mole of ionic solid
Usually reported as exothermic process
Energy required to separate 1 mole of ionic solid into gaseous ions

5. Lattice Energy
Depends directly on magnitude of charges and inversely on size of ions
Larger charges means more coulombic attraction
Larger ion size means less coulombic attraction

6. Ionic Compounds Metal and nonmetal High Melting and Boiling Points
Hard Crystalline solids
Brittle
Break when struck (Cleave)
Soluble in water
Conduct electricity when melted or dissolved
Do not conduct as solids

7. Covalent Bonding
Sharing of valence electrons due to equal or near equal electronegativity of 2 nonmetals

8. Bond Enthalpy Amount of energy required to break a bond
Amount of energy released when a bond forms
Larger energy = stronger bond
ΔH = Σreactants - Σproducts

9. Bond Polarity
Bond type is based on difference in electronegativity between bonding atoms
Larger difference = more polar
Bond polarity is a continuous spectrum
Nonpolar
Covalent
Polar
Covalent
Ionic
ΔEN

10. Dipole Moment
Measure of Bond Polarity

11. Molecular Polarity

12. Molecular Polarity

14. F Lewis Structures Electron Dot Diagrams Rules for individual atoms
Shows valence electrons only
Rules for individual atoms
Start on any side
First two get paired together
Next three are separated
Fill in as needed F

15. F- Na+ Lewis Structures Monatomic Ions
Positive ions tend to have 0 dots
Lost all valence electrons
Negative ions tend to have 8 dots
Gained enough to fill valence shell
Can have brackets around ion with charge outside brackets
Na+ F-

16. Lewis Structures
Ionic Compounds
NaCl Cl
[ ] -
Na+

17. Lewis Structures Covalent Compounds Line for bonding electrons
1 line = single bond
2 lines = double bond
3 lines = triple bond
Dots for nonbonding electrons

18. Lewis Structures Determine number valence electrons
Arrange atoms around a central atom with single bonds
Usually least electronegative element
excluding H
Fill in octet for outer atoms
Fill in octet for central atom
If not enough to fill central atom’s octet, make multiple bonds to central atom

19. Example
CO2
16 valence electrons O C O

20. Lewis Structures
Polyatomic Ions O N O O

21. Lewis Structures Exceptions to Octet Rule More than octet
Less than octet
Odd number of electrons

22. Lewis Structures
SF6 F F F S F F F

23. Lewis Structures
BF3 F B F F

24. Lewis Structures NO N O

26. Formal Charge
Calculated quantity used to determine the validity of Lewis structures
For each atom,
# of valence electrons – (nonbonding e- + ½ bonding e-) C N H

27. Formal Charge Best Lewis Structures
Have fewest and/or smallest charges
Negative charge on most electronegative element

28. Formal Charge Example -1 S O H +2 S O H -1

29. Lewis Structures
Benzene, C6H6 H C

30. Lewis Structures
Benzene, C6H6 H H C C H C C H C C H H

31. Resonance Two or more equally valid, but different Lewis Structures
Differ only in the position of the electrons
Use double headed arrow between resonance structures

32. Resonance for Benzene Either structure would suggest
Double bonds should be shorter than single bonds
Electrons are localized between carbon atoms

33. Resonance for Benzene Experimental evidence shows
All bonds are equal in length
Length halfway between single and double bond
Electrons are “delocalized”
Spread around ring equally

35. VSEPR Theory Valence Shell Electron Pair Repulsion
Model accounts for the shape of the molecule
Lewis Structures don’t always show proper shape
Electron groups repel each other and try to spread as far apart as possible
Electron groups (domains) are lone pairs, single bonds, double bonds, and triple bonds

36. Electron Geometry
There are 5 different possible arrangements around a central atom
2 electron groups up to 6 electron groups

37. Electron Geometry 2 electron groups around a central atom
Electron groups on opposite sides
Bond Angle of 180°
Linear Shape

38. Electron Geometry 3 electron groups around a central atom
Bond Angle of 120°
Trigonal Planar

39. Electron Geometry 4 electron groups around a central atom Tetrahedral
Bond Angle of 109.5°
Tetrahedral

40. Electron Geometry 5 electron groups around a central atom
3 equatorial, 2 axial
Equatorial angle of 120°, axial angle of 90°
Trigonal bipyramidal

41. Electron Geometry 6 electron groups around a central atom Octahedral
Bond angle of 90°
Octahedral

42. Molecular Geometry
Molecular Geometry is defined by the position of the bonded atoms in the molecule, not the lone pairs
All binary molecules are linear
Lone pairs still impact the shape due to their repulsion
Same electron geometries may have different molecular geometries

43. Molecular Geometry Bent 3 electron groups, 2 bonding & 1 lone pair
SO2
4 electron groups, 2 bonding & 2 lone pairs
H2O

44. Molecular Geometry Trigonal Pyramidal
4 electron groups, 3 bonding & 1 lone pair
NH3

45. Molecular Geometry Seesaw 5 electron groups, 4 bonding & 1 lone pair
SF4

46. Molecular Geometry T-shaped
5 electron groups, 3 bonding & 2 lone pairs
BrF3

47. Molecular Geometry Linear 5 electron groups, 2 bonding & 3 lone pairs
XeF2

48. Molecular Geometry Square Pyramidal
6 electron groups, 5 bonding & 1 lone pair
BrF5

49. Molecular Geometry Square Planar
6 electron groups, 4 bonding & 2 lone pairs
XeF4

50. Bond Angles Angle between 2 atoms bonded to a central atom
Other bonds and/or lone pairs in the molecule affect bond angle

51. Lone Pairs
Electrons in a lone pair take up more space than bonded atoms due to repulsion between the electrons
Multiple bonds also take up more space

52. Larger Molecules
For larger molecules, we talk about the geometry around individual atoms

54. Valence Bond Theory
Orbitals must overlap in order for atoms to share electrons and bond together

55. Internuclear Distance
Balancing of attraction between nuclei and electrons with repulsion between electrons
Nuclei repel each other if atoms get too close together

56. Internuclear Distance

57. Bonding Carbon has an electron configuration of 1s22s22p2
Carbon only has 2 unpaired electrons
Carbon should only make 2 bonds

58. Hybrid Orbitals Orbitals can hybridize to form new equal orbitals
Able to form more bonds
Carbon mixes 2s and all 2p orbitals to make 4 equal sp3 hybrid orbitals

59. Hybrid Orbitals

60. Hybrid Orbitals

61. Hybrid Orbitals

62. Sigma and Pi Bonds
There are 2 ways that orbitals can overlap to form bonds
Sigma bonds
Pi bonds

63. Sigma (σ) Bonds
Head to head overlap
All single bonds are sigma bonds

64. Pi (π) Bonds Side to side overlap
Multiple bonds have both sigma and pi bonds
Double bonds
1 sigma, 1 pi
Triple Bonds
1 sigma, 2 pi

65. Multiple Bonds

66. Multiple Bonds

67. Resonance

68. Resonance

70. Metallic Bonding
Bonding in metals is due to highly mobile valence electrons
Delocalized electrons
Metal nuclei and inner (core) electrons act as a cation, keeping the valence electrons within the sample
Ability of the electrons to move throughout the entire sample give metals their unique properties

71. Metallic Bonding Metallic Properties
Conductivity of heat and electricity
Malleable
Ductile

72. Alloys Metallic material that is composed of more than one element
More rigid, less malleable and ductile than pure metal
Two types
Substitutional
Interstitial

73. Substitutional Alloy
One metal atom takes the place of original metal atom
Brass, Bronze

74. Interstitial Alloys
Small, often nonmetallic, atoms fit in between the metal atoms in a crystal
Steel