1. CHEMICAL BONDING Cocaine
To play the movies and simulations included, view the presentation in Slide Show Mode.

2. Chemical Bonding Problems and questions —
How is a molecule or polyatomic ion held together?
Why are atoms distributed at strange angles?
Why are molecules not flat?
Can we predict the structure?
How is structure related to chemical and physical properties?

3. Forms of Chemical Bonds
There are 2 extreme forms of connecting or bonding atoms:
Ionic—complete transfer of 1 or more electrons from one atom to another
Covalent—some valence electrons shared between atoms
Most bonds are somewhere in between.

4. Ionic Bonds
Essentially complete electron transfer from an element of low IE (metal) to an element of high affinity for electrons (nonmetal)
2 Na(s) + Cl2(g) --->
2 Na Cl-
Therefore, ionic compds. exist primarily between metals at left of periodic table (Grps 1A and 2A and transition metals) and nonmetals at right (O and halogens).

5. Covalent Bonding
The bond arises from the mutual attraction of 2 nuclei for the same electrons. Electron sharing results. (Screen 9.5)
Bond is a balance of attractive and repulsive forces.

6. Chemical Bonding: Objectives
Objectives are to understand:
1. valence e- distribution in molecules and ions.
2. molecular structures
3. bond properties and their effect on molecular properties.

7. Electron Distribution in Molecules
Electron distribution is depicted with Lewis electron dot structures
Valence electrons are distributed as shared or BOND PAIRS and unshared or LONE PAIRS.
G. N. Lewis

8. Bond and Lone Pairs
Valence electrons are distributed as shared or BOND PAIRS and unshared or LONE PAIRS. • •• H Cl
shared or
bond pair
lone pair (LP)
This is called a LEWIS ELECTRON DOT structure.

9. Note that each atom has a single, unpaired electron.
Bond Formation
A bond can result from a “head-to-head” overlap of atomic orbitals on neighboring atoms. Cl H •• • +
Overlap of H (1s) and Cl (2p)
Note that each atom has a single, unpaired electron.

10. Valence Electrons
Electrons are divided between core and valence electrons
B 1s2 2s2 2p1
Core = [He] , valence = 2s2 2p1
Br [Ar] 3d10 4s2 4p5
Core = [Ar] 3d10 , valence = 4s2 4p5

11. Rules of the Game
No. of valence electrons of a main group atom = Group number
•For Groups 1A-4A (14), no. of bond pairs = group number.
• For Groups 5A (15)-7A (17), BP’s = 8 - Grp. No.

12. Rules of the Game •No. of valence electrons of an atom = Group number
•For Groups 1A-4A (14), no. of bond pairs = group number
• For Groups 5A (15)-7A (17), BP’s = 8 - Grp. No.
•Except for H (and sometimes atoms of 3rd and higher periods),
BP’s + LP’s = 4
This observation is called the
OCTET RULE

13. Building a Dot Structure
Ammonia, NH3
1. Decide on the central atom; never H.
Central atom is atom of lowest affinity for electrons.
Therefore, N is central
2. Count valence electrons
H = 1 and N = 5
Total = (3 x 1) + 5
= 8 electrons / 4 pairs

14. Building a Dot Structure
3. Form a single bond between the central atom and each surrounding atom H N
4. Remaining electrons form LONE PAIRS to complete octet as needed. H •• N
3 BOND PAIRS and 1 LONE PAIR.
Note that N has a share in 4 pairs (8 electrons), while H shares 1 pair.

15. Sulfite ion, SO32- Step 1. Central atom = S
Step 2. Count valence electrons S = 6
3 x O = 3 x 6 = 18
Negative charge = 2
TOTAL = 26 e- or 13 pairs
Step 3. Form bonds
10 pairs of electrons are
now left.

16. Sulfite ion, SO32-
Remaining pairs become lone pairs, first on outside atoms and then on central atom. • •• O S ••
Each atom is surrounded by an octet of electrons.

17. Carbon Dioxide, CO2 1. Central atom = _______
2. Valence electrons = __ or __ pairs
3. Form bonds.
This leaves 6 pairs.
4. Place lone pairs on outer atoms.

18. Carbon Dioxide, CO2 4. Place lone pairs on outer atoms.
5. So that C has an octet, we shall form DOUBLE BONDS between C and O.
The second bonding pair forms a pi (π) bond.

19. Double and even triple bonds are commonly observed for C, N, P, O, and S
H2CO
SO3
C2F4

20. Sulfur Dioxide, SO2 O S 1. Central atom = S
2. Valence electrons = 18 or 9 pairs
3. Form double bond so that S has an octet — but note that there are two ways of doing this.
bring in
left pair
OR bring in
right pair • O S ••

21. Sulfur Dioxide, SO2 This leads to the following structures.
These equivalent structures are called
RESONANCE STRUCTURES. The true electronic structure is a HYBRID of the two.

22. Urea, (NH2)2CO

23. Urea, (NH2)2CO 1. Number of valence electrons = 24 e-
2. Draw sigma bonds.

24. Urea, (NH2)2CO
3. Place remaining electron pairs in the molecule.

25. Urea, (NH2)2CO
4. Complete C atom octet with double bond.

26. Violations of the Octet Rule
Usually occurs with B and elements of higher periods.
SF4
BF3

27. Boron Trifluoride Central atom = _____________
Valence electrons = __________ or electron pairs = __________
Assemble dot structure
The B atom has a share in only 6 pairs of electrons (or 3 pairs). B atom in many molecules is electron deficient.

28. Sulfur Tetrafluoride, SF4
Central atom =
Valence electrons = ___ or ___ pairs.
Form sigma bonds and distribute electron pairs.
5 pairs around the S atom. A common occurrence outside the 2nd period.

29. Formal Atom Charges
Atoms in molecules often bear a charge (+ or -).
The predominant resonance structure of a molecule is the one with charges as close to 0 as possible.
Formal charge = Group no. – 1/2 (no. of BE) (no. of LP electrons)

30. Carbon Dioxide, CO2 6 - ( 1 / 2 ) 4 = • O C 4 - ( 1 / 2 ) 8 =

31. Calculated Partial Charges in CO2
Yellow = negative & red = positive
Relative size = relative charge

32. Thiocyanate Ion, SCN- S N C 6 - (1/2)(2) - 6 = -1 5 - (1/2)(6) - 2 = 0• S N C
4 - (1/2)(8) - 0 = 0

33. Which is the most important resonance form?
Thiocyanate Ion, SCN- • S N C • S N C • S N C
Which is the most important resonance form?

34. Calculated Partial Charges in SCN-
All atoms negative, but most on the S • S N C

35. Boron Trifluoride, BF3 F B•• • B +1 -1
What if we form a B—F double bond to satisfy the B atom octet?

36. Is There a B=F Double Bond in BF3
Calc’d partial charges in BF3
F is negative and B is positive

37. MOLECULAR GEOMETRY

38. VSEPR MOLECULAR GEOMETRY Valence Shell Electron Pair Repulsion theory.
Molecule adopts the shape that minimizes the electron pair repulsions.
VSEPR
Valence Shell Electron Pair Repulsion theory.
Most important factor in determining geometry is relative repulsion between electron pairs.

39. Electron Pair Geometries Figure 9.12

44. Electron Pair Geometries Figure 9.12

45. Structure Determination by VSEPRH •• N
Ammonia, NH3
1. Draw electron dot structure
2. Count BP’s and LP’s = 4
3. The 4 electron pairs are at the corners of a tetrahedron.

46. Structure Determination by VSEPR
Ammonia, NH3
There are 4 electron pairs at the corners of a tetrahedron.
The ELECTRON PAIR GEOMETRY is tetrahedral.

47. Structure Determination by VSEPR
Ammonia, NH3
The electron pair geometry is tetrahedral.
The MOLECULAR GEOMETRY — the positions of the atoms — is PYRAMIDAL.

48. Structure Determination by VSEPR
Water, H2O
1. Draw electron dot structure
2. Count BP’s and LP’s = 4
3. The 4 electron pairs are at the corners of a tetrahedron.
The electron pair geometry is TETRAHEDRAL.

49. Structure Determination by VSEPR
Water, H2O
The electron pair geometry is TETRAHEDRAL
The molecular geometry is BENT.

50. Geometries for Four Electron Pairs Figure 9.13

51. Structure Determination by VSEPR• C H O
Formaldehyde, CH2O
1. Draw electron dot structure
2. Count BP’s and LP’s at C
3. There are 3 electron “lumps” around C at the corners of a planar triangle. • C H O
The electron pair geometry is PLANAR TRIGONAL with 120o bond angles.

52. Structure Determination by VSEPR
Formaldehyde, CH2O
The electron pair geometry is PLANAR TRIGONAL
The molecular geometry is also planar trigonal.

53. Structure Determination by VSEPR
Methanol, CH3OH H ••
H—C—O—H
Define H-C-H and C-O-H bond angles
109˚
109˚
H-C-H = 109o
C-O-H = 109o
In both cases the atom is surrounded by 4 electron pairs.

54. Structure Determination by VSEPR
Acetonitrile, CH3CN •• H
H—C—C N
Define unique bond angles
H-C-H = 109o
C-C-N = 180o
109˚
180˚
One C is surrounded by 4 electron “lumps” and the other by 2 “lumps”

55. Phenylalanine, an amino acid

56. Phenylalanine

57. Structures with Central Atoms with More Than or Less Than 4 Electron Pairs
Often occurs with Group 3A elements and with those of 3rd period and higher.

58. Boron Compounds Consider boron trifluoride, BF3
The B atom is surrounded by only 3 electron pairs.
Bond angles are 120o
Geometry described as planar trigonal

59. Compounds with 5 or 6 Pairs Around the Central Atom
5 electron pairs

60. Sulfur Tetrafluoride, SF4
Number of valence electrons = 34
Central atom = S
Dot structure
Electron pair geometry
--> trigonal bipyramid (because there are 5 pairs around the S)

61. Sulfur Tetrafluoride, SF4
Lone pair is in the equator because it requires more room.

62. Molecular Geometries for Five Electron Pairs Figure 9.14

63. Compounds with 5 or 6 Pairs Around the Central Atom
6 electron pairs

64. Molecular Geometries for Six Electron Pairs Figure 9.14

65. Bond Properties
What is the effect of bonding and structure on molecular properties?
Buckyball in HIV-protease

66. # of bonds between a pair of atoms
Bond Order
# of bonds between a pair of atoms
Double bond
Single bond
Acrylonitrile
Triple bond

67. Bond Order The N—O bond order = 1.5
Fractional bond orders occur in molecules with resonance structures.
Consider NO2-
The N—O bond order = 1.5

68. Bond Order 414 kJ 123 pm 745 kJ 110 pm (a) bond strength
Bond order is proportional to two important bond properties:
(a) bond strength
(b) bond length
745 kJ
414 kJ
110 pm
123 pm

69. Bond Length
Bond length is the distance between the nuclei of two bonded atoms.

70. Bond Length H—F H—Cl H—I Bond length depends on size of bonded atoms.
Bond distances measured using CAChe software. In Angstrom units where 1 A = 10-2 pm.

71. Bond Length Bond length depends on bond order.
Bond distances measured using CAChe software. In Angstrom units where 1 A = 10-2 pm.

72. Bond Strength
—measured by the energy req’d to break a bond. See Table 9.9

73. Bond Strength
—measured by the energy req’d to break a bond. See Table 9.9.
BOND STRENGTH (kJ/mol)
H—H 436
C—C 346
C=C 602
CC
NN 945
The GREATER the number of bonds (bond order) the HIGHER the bond strength and the SHORTER the bond.

74. Bond Strength 1 142 pm 210 kJ/mol 2 121 494 1.5 128 ?
Bond Order Length Strength
HO—OH
O=O 1
142 pm
210 kJ/mol 2
121
494
1.5
128 ?

75. Using Bond Energies H—H = 436 kJ/mol Cl—Cl = 242 kJ/mol
Estimate the energy of the reaction
H—H + Cl—Cl ----> 2 H—Cl
Net energy = ∆Hrxn =
= energy required to break bonds
- energy evolved when bonds are made
H—H = 436 kJ/mol
Cl—Cl = 242 kJ/mol
H—Cl = 432 kJ/mol

76. Sum of H-H + Cl-Cl bond energies = 436 kJ + 242 kJ = +678 kJ
Using Bond Energies
Estimate the energy of the reaction
H—H + Cl—Cl ----> 2 H—Cl
H—H = 436 kJ/mol
Cl—Cl = 242 kJ/mol
H—Cl = 432 kJ/mol
Sum of H-H + Cl-Cl bond energies = 436 kJ kJ = +678 kJ
2 mol H-Cl bond energies = 864 kJ
Net = ∆H = kJ kJ = -186 kJ

77. Using Bond Energies Estimate the energy of the reaction
2 H—O—O—H ----> O=O + 2 H—O—H
Is the reaction exo- or endothermic?
Which is larger:
A) energy req’d to break bonds
B) or energy evolved on making bonds?

78. Using Bond Energies TOTAL ENERGY to break bonds = 2144 kJ
2 H—O—O—H > O=O + 2 H—O—H
Energy required to break bonds:
break 4 mol of O—H bonds = 4 (463 kJ)
break 2 mol O—O bonds = 2 (146 kJ)
TOTAL ENERGY to break bonds = 2144 kJ
TOTAL ENERGY evolved on making O=O bonds and 4 O-H bonds bonds = kJ

79. Using Bond Energies Net energy = +2144 kJ - 2350 kJ = - 206 kJ
2 H—O—O—H > O=O + 2 H—O—H
Net energy = kJ kJ = kJ
The reaction is exothermic!
More energy is evolved on making bonds than is expended in breaking bonds.

80. Molecular Polarity Boiling point = 100 ˚C Boiling point = -161 ˚C
Why do water and methane differ so much in their boiling points?
Why do ionic compounds dissolve in water?

81. Why do CuCl2, acetic acid, and ethanol dissolve in water?

82. Bond Polarity
HCl is POLAR because it has a positive end and a negative end.
Cl has a greater share in bonding electrons than does H.
Cl has slight negative charge (-d) and H has slight positive charge (+ d)

83. Bond Polarity
This model, calc’d using CAChe software, shows that H is + (red) and Cl is - (yellow). Calc’d charge is + or (On CD see PARTCHRG folder in MODELS.)

84. Bond Polarity
Due to the bond polarity, the H—Cl bond energy is GREATER than expected for a “pure” covalent bond.
BOND ENERGY
“pure” bond 339 kJ/mol calc’d
real bond 432 kJ/mol measured
Difference = 92 kJ. This difference is proportional to the difference in ELECTRONEGATIVITY, .

85. Electronegativity, 
 is a measure of the ability of an atom in a molecule to attract electrons to itself.
Concept proposed by
Linus Pauling

86. Linus Pauling,
The only person to receive two unshared Nobel prizes (for Peace and Chemistry).
Chemistry areas: bonding, electronegativity, protein structure

87. Electronegativity Figure 9.9

88. Electronegativity,  F has maximum .
See Figure 9.9
F has maximum .
Atom with lowest  is the center atom in most molecules.
Relative values of  determine BOND POLARITY (and point of attack on a molecule).

89. Bond Polarity Which bond is more polar (or DIPOLAR)? O—H O—F 
OH is more polar than OF
and polarity is “reversed.”

90. Molecular Polarity
Molecules—such as HCl and H2O— can be POLAR (or dipolar).
They have a DIPOLE MOMENT. The polar HCl molecule will turn to align with an electric field.
Figure 9.15

91. Molecular Polarity
The magnitude of the dipole is given in Debye units. Named for Peter Debye ( ). Rec’d 1936 Nobel prize for work on x-ray diffraction and dipole moments.

92. Why are some molecules polar but others are not?
Dipole Moments
Why are some molecules polar but others are not?

93. Molecular Polarity All above are NOT polar Molecules will be polar if
a) bonds are polar
AND
b) the molecule is NOT “symmetric”
All above are NOT polar

94. Compare CO2 and H2O. Which one is polar?
Polar or Nonpolar?
Compare CO2 and H2O. Which one is polar?

95. Carbon Dioxide
CO2 is NOT polar even though the CO bonds are polar.
CO2 is symmetrical.
+1.5
-0.75
Positive C atom is reason CO2 + H2O gives H2CO3

96. Consequences of H2O Polarity
Microwave oven
Consequences of H2O Polarity

97. Polar or Nonpolar?
Consider AB3 molecules: BF3, Cl2CO, and NH3.

98. B—F bonds in BF3 are polar. But molecule is symmetrical and NOT polar
Molecular Polarity, BF3
B atom is positive and F atoms are negative.
B—F bonds in BF3 are polar.
But molecule is symmetrical and NOT polar

99. Molecular Polarity, HBF2
B atom is positive but H & F atoms are negative.
B—F and B—H bonds in HBF2 are polar. But molecule is NOT symmetrical and is polar.

100. Is Methane, CH4, Polar?
Methane is symmetrical and is NOT polar.

101. Is CH3F Polar?
C—F bond is very polar. Molecule is not symmetrical and so is polar.

102. Substituted Ethylene C—F bonds are MUCH more polar than C—H bonds.
Because both C—F bonds are on same side of molecule, molecule is POLAR.

103. Substituted Ethylene C—F bonds are MUCH more polar than C—H bonds.
Because both C—F bonds are on opposing ends of molecule, molecule is NOT POLAR.