1. AP Notes Chapter 8
Bonding and Molecular Structure: Fundamental Concepts
Valence e- and Bonding
Covalent
Ionic
Resonance & Exceptions to Octet Rule
Bond Energy & Length
Structure, Shape & Polarity of Compounds

2. What is a Bond? A force that holds atoms together. Why?
We will look at it in terms of energy.
Bond energy the energy required to break a bond.
Why are compounds formed?
Because it gives the system the lowest energy.

3. Covalent compounds?
The electrons in each atom are attracted to the nucleus of the other.
The electrons repel each other,
The nuclei repel each other.
The reach a distance with the lowest possible energy.
The distance between is the bond length.

4. Thus Hydrogen is Diatomic!
Bond Formation

5. Covalent Character e-

6. Why Isn’t Helium Diatomic?. . E
He + He
He2
Inter-nuclear Distance . .

7. 2p ____ ____ ___ ___ ____ ____ 2p 2s ____ ____ 2s
F + F F2
2p ____ ____ ___ ___ ____ ____ 2p 2s ____ ____ 2s
F F

8. Ionic Bonding
An atom with a low ionization energy reacts with an atom with high electron affinity.
The electron moves.
Opposite charges hold the atoms together.

9. Li + Cl 1s22s1 [Ne] 3s23p5 2s ___ 3p _____ _____ ___ 1s _____ 3s _____ [Ne]

10. Li + Cl 2s ___ 3P _____ _____ _____ 1s _____ 3s _____ [Ne]

11. LiCl 2s ___ 3P _____ _____ _____ 1s _____ 3s _____ [Ne]

12. Electronegativity
Describes the relative ability of an atom within a molecule to attract a shared pair of electrons to itself.

13. Pauling electronegativity values, which are unit-less, are the norm.

14. Electronegativity Range from 0.7 to 4.0
Figure 9.9 – Kotz & Treichel

15. Bond: A - B
DEN = | ENA - ENB |

16. Bond Character “Ionic Bond” - Principally Ionic Character
“Covalent Bond” - Principally Covalent Character

17. Determining Principal Character of Bond
covalent
ionic
EN ~0 ~4
1.7

18. F - F EN = 0
Non-polar

19. N - O EN = | | = 0.5 O N
Slightly polar

21. Ionic Bond with some covalent character
Ca - O  EN = | | = 2.5 Ca O
Ionic Bond with some covalent character

23. Electronegativity
The ability of an electron to attract shared electrons to itself.
Pauling method
Imaginary molecule HX
Expected H-X energy = H-H energy + X-X energy
D = (H-X) actual - (H-X)expected

24. Electronegativity D is known for almost every element
Gives us relative electronegativities of all elements.
Tends to increase left to right.
decreases as you go down a group.
Noble gases aren’t discussed.
Difference in electronegativity between atoms tells us how polar.

25. Electronegativity
difference
Bond
Type
Zero
Covalent
Covalent Character
decreases
Ionic Character increases
Polar
Covalent
Intermediate
Ionic
Large

26. Dipole Moments
A molecule with a center of negative charge and a center of positive charge is dipolar (two poles),
or has a dipole moment.
Center of charge doesn’t have to be on an atom.
Will line up in the presence of an electric field.

27. How It is drawn
H - F d+ d-

28. Which Molecules Have Them?
Any two atom molecule with a polar bond.
With three or more atoms there are two considerations.
There must be a polar bond.
Geometry can’t cancel it out.

31. Ionic Radii -- Cations

32. Ionic Radii -- Anions

34. Vector Sum of Bond Polarities
Molecular Polarity
Vector Sum of Bond Polarities

35. Covalent BOND w/much ionic character, BUT NON-POLAR molecule
MgBr2
Mg - Br EN = | | = 1.6 Mg Br Br
Covalent BOND w/much ionic character, BUT NON-POLAR molecule

39. Lewis
Structures

40. The most important requirement for the formation of a stable compound is that the atoms achieve noble gas e- configuration

41. Valence Shell Electron Pair Repulsion Model (VSEPR)
The structure around a given atom is determined principally by minimizing electron-pair repulsions

42. VSEPR Electron pairs Bond Angles Underlying Shape 2 180° Linear 3 120°
Trigonal Planar 4
109.5°
Tetrahedral 5
90° &
120°
Trigonal Bipyramidal 6
90°
Octagonal

43. LEWIS STRUCTURES : draw skeleton of species : count e- in species
: subtract 2 e- for each bond in skeleton
: distribute remaining e-

44. Distinguish Between
ELECTRONIC
Geometry
&
MOLECULAR

45. CH4 Bond angle = 109.50 Electronic geometry: tetrahedral
Molecular geometry: tetrahedral

46. H3O+ Bond angle ~ 1070 Electronic geometry: tetrahedral
Molecular geometry: trigonal pyramidal

47. H2O Bond angle ~ 104.50 Electronic geometry: tetrahedral
Molecular geometry: bent

48. NH2- Bond angle ~ 104.50 Electronic geometry: tetrahedral
Molecular geometry: bent

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

51. BaI2 Bond angle =1800 Electronic geometry: linear
Molecular geometry: linear

52. BF3 Bond angle =1200 Electronic geometry: trigonal planar
Molecular geometry: trigonal planar

53. PF5 Bond angle = 1200 & 900 Electronic geometry: trigonal bipyramidal
Molecular geometry: trigonal bipyramidal

54. SF4 Bond angle = 1200 & 900 Electronic geometry: trigonal bipyramidal
Molecular geometry: see-saw

55. ICl3 Bond angle <= 900 Electronic geometry: trigonal bipyramidal
Molecular geometry: T-shape

56. I3- Bond angle = 1800 Electronic geometry: trigonal bipyramid
Molecular geometry: linear

58. PCl6- Bond angle = 900 Electronic geometry: octahedral
Molecular geometry: octahedral

59. BrF5 Bond angle ~ 900 Electronic geometry: octahedral
Molecular geometry: square pyramidal

60. ICl4- Bond angle = 900 Electronic geometry: octahedral
Molecular geometry: square planar

61. Actual shape Non-BondingPairs ElectronPairs BondingPairs Shape 2 2
linear 3 3
trigonal planar 3 2 1
bent 4 4
tetrahedral 4 3 1
trigonal pyramidal 4 2 2
bent

62. Actual Shape Non-BondingPairs ElectronPairs BondingPairs Shape 5 5
trigonal bipyrimidal 5 4 1
See-saw 5 3 2
T-shaped 5 2 3
linear

63. Actual Shape Non-BondingPairs ElectronPairs BondingPairs Shape 6 6
Octahedral 6 5 1
Square Pyramidal 6 4 2
Square Planar 6 3 3
T-shaped 6 2 1
linear

64. What happens when there are not enough electrons to “satisfy” the central atom?

65. EXAMPLES
Ethene
Acetic Acid
Oxygen
Nitrogen

66. RESONANCE
&
FORMAL CHARGE

67. Resonance
Sometimes there is more than one valid structure for an molecule or ion.
NO3-
Use double arrows to indicate it is the “average” of the structures.
It doesn’t switch between them.
NO2-
Localized electron model is based on pairs of electrons, doesn’t deal with odd numbers.

68. EXAMPLES
Nitrate ion
Ozone

69. FORMAL CHARGE
the charge assigned to an atom in a molecule or polyatomic ion
FC atom = Family# - [LPE + ½(BE)]
Sum FC’s atoms = ion charge
Closer sum FC’s is to zero more stable

70. Formal Charge
For molecules and polyatomic ions that exceed the octet there are several different structures.
Use charges on atoms to help decide which.
Trying to use the oxidation numbers to put charges on atoms in molecules doesn’t work.

71. Formal Charge
The difference between the number of valence electrons on the free atom and that assigned in the molecule.
We count half the electrons in each bond as “belonging” to the atom.
SO4-2
Molecules try to achieve as low a formal charge as possible.
Negative formal charges should be on electronegative elements.

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

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

74. A useful equation (happy-have) / 2 = bonds POCl3 P is central atom
SO S is central atom
SO3-2 S is central atom
PO P is central atom
SCl2 S is central atom

75. Exceptions to the octet
BH3
Be and B often do not achieve octet
Have less than and octet, for electron deficient molecules.
SF6
Third row and larger elements can exceed the octet
Use 3d orbitals?
I3-

76. Exceptions to the octet
When we must exceed the octet, extra electrons go on central atom.
ClF3
XeO3
ICl4-
BeCl2

77. If nonequivalent Lewis structures exist, the one(s) that best describe the bonding in the species has...

78. FAVORED LEWIS STRUCTURES
1. formal charges closest to zero
2. negative formal charge
is on the most electronegative atom

79. EXAMPLES
Carbon dioxide
Thiocyanate ion
Sulfate ion

80. BOND
ENERGY & LENGTH

81. Bond Energies E = (Bonds Broken) – (Bonds Made)

82. Bonds form between atoms because bonded atoms exhibit a lower energy.
Thus, energy is required to break bonds and energy is released when bonds are formed.

83. Bond Order = # bonds to a specific set of elements C-C the BO=1 C=C the BO=2 C C the BO=3 Fractions are possible

84. COVALENT BONDS
Bond Dissociation
Energy
Table 9.9 (text)

85. Bond Energy
(kJ/mol)
H-F
H-Cl
H-Br
H-I

86. Bond Energy
(kJ/mol)
Cl-Cl
Br-Br 193
I-I

87. Bond Energy
(kJ/mol)

88. Bond Energy
(kJ/mol)

89. Use bond energies to predict Hc for acetylene (C2H2).

90. Energy
Internuclear Distance

91. Energy
Internuclear Distance

92. Energy
Internuclear Distance

93. Energy
Internuclear Distance

94. Energy
Bond Length
Internuclear Distance

95. Energy
Bond Energy
Internuclear Distance

96. Bond: Energy Length
(kJ/mol) (pm)

97. Bond: Energy Length
(kJ/mol) (pm)

98. Binary Ionic Compounds
metal(s) + non-metal (g) ---> salt(s)

99. M+(g) + NM-(g) --> M-NM
Lattice Energy
Energy change occurring when separated gaseous ions are packed together to form an ionic solid
M+(g) + NM-(g) --> M-NM

100. What is the lattice energy of NaCl(s)? Na+(g) + Cl-(g) ---> NaCl(s)

101. Lattice Energies LiCl 834 : BeCl2 3004 NaCl 769 : MgCl2 2326
KCl : CaCl
Li2O : BeO
Na2O : MgO
K2O : CaO

102. LE = Lattice Energy
Where: k = proportionality constant dependent on structure of solid and on electron configuration of the ions
Where: Q1 & Q2 = charges on the ions
Where: r = the shortest distance
between the centers of the cation and anion