1. Bonding and Molecular Structure: Fundamental Concepts
Chapter 9
Bonding and Molecular Structure:
Fundamental Concepts
Dr. S. M. Condren

2. Chemical Bonding How is a molecule or polyatomic ion held together?
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?
Dr. S. M. Condren

3. Structure & Bonding NN triple bond. Molecule is unreactive
White phosphorus is a tetrahedron of P atoms.
Very reactive!
Red phosphorus, a polymer. Used in matches.
Less reactive!
Dr. S. M. Condren

4. 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.
Dr. S. M. Condren

5. Ionic Compounds 2 Na(s) + Cl2(g) ---> 2 Na+ + 2 Cl- Metal of low IE
Nonmetal of high EA
2 Na(s) + Cl2(g) ---> 2 Na Cl-
Dr. S. M. Condren

6. Covalent Bonding
The bond arises from the mutual attraction of 2 nuclei for the same electrons. Electron sharing results.
Bond is a balance of attractive and repulsive forces.
Dr. S. M. Condren

7. 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 •• • +
Note that each atom has a single, unpaired electron.
Overlap of H (1s) and Cl (3p)
Dr. S. M. Condren

8. 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.
Dr. S. M. Condren

9. 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
Dr. S. M. Condren

10. Bond and Lone Pairs
Valence electrons are distributed as shared or BOND PAIRS and unshared or LONE PAIRS. • •• H Cl
lone pairs
shared or
bond pair
This is called a LEWIS ELECTRON DOT structure.
Dr. S. M. Condren

11. 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
Dr. S. M. Condren

12. Rules of the Game
No. of valence electrons of a main group atom = Group number
•For Groups 1A-4A, no. of bond pairs = group number.
• For Groups 5A -7A, BP’s = 8 - Grp. No.
Group 3A
Group 5A
Dr. S. M. Condren

13. Rules of the Game No. of valence electrons of an atom = Group number
For Groups 1A-4A, no. of bond pairs = group number
For Groups 5A -7A, 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
Dr. S. M. Condren

14. Hydrophobic vs. Hydrophilic
Hydrophobic - translation of Greek – water fear
Hydrophilic – translation of Greek – water friendship
Dr. S. M. Condren

15. 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
Dr. S. M. Condren

16. 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.
Dr. S. M. Condren

17. Lewis Structures CH4 methane Step 1. Central atom = C
Step 2. Count valence electrons
C = 4
4 x H = 4 x 1 = 4
TOTAL = 8 e- or 4 pairs
Step 3. Form bonds
Dr. S. M. Condren

18. Lewis Structures
C2H6 ethane
Dr. S. M. Condren

19. Multiple Covalent Bonds
double bond => 2 pairs shared
triple bond => 3 pairs shared
normally occurs between:
C atoms; N atoms; O atoms;
a C atom and a N, O or S atom
a N atom and a O or S atom
a S atom and an O atom
Dr. S. M. Condren

20. Double and even triple bonds are commonly observed for C, N, P, O, and S
H2CO
SO3
C2F4
Dr. S. M. Condren

21. Carbon Dioxide, CO2 C 4 2 1. Central atom = _______
2. Valence electrons = __ or __ pairs
3. Form bonds.
This leaves 6 pairs.
Dr. S. M. Condren

22. 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.
Dr. S. M. Condren

23. Steps to form Lewis Electron Dot Structure
1. Central atom = _______
2. Valence electrons = __ or __ pairs
Form bonds.
Place lone pairs on outer atoms.
Form multiple bonds as necessary to obey Lewis’ “Octet Rule”.
Remember that there are MANY compounds that do not obey the Octet Rule.
Dr. S. M. Condren

24. Lewis Structures
CO carbon monoxide
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25. Exceptions to Octet Rule
NO nitric oxide
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26. Exceptions to Octet Rule
NO2 nitrogen dioxide
resonance
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27. Exceptions to Octet Rule
PF5
expanded octet
Dr. S. M. Condren

28. Exceptions to Octet Rule
expanded octet
SF4
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29. Exceptions to Octet Rule
SF expanded octet
Dr. S. M. Condren

30. 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 number – 1/2 (no. of bonding electrons) - (no. of LP electrons)
Dr. S. M. Condren

31. Carbon Dioxide, CO2 O C +6 - ( 1 / 2 ) 4 = +4 - ( 1 / 2 ) 8 = •• O C +4 - ( 1 / 2 ) 8 =
Dr. S. M. Condren

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
Dr. S. M. Condren

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?
Dr. S. M. Condren

34. MOLECULAR GEOMETRY
Dr. S. M. Condren

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

36. Electron Pair Geometries
Dr. S. M. Condren

37. 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 ••
Dr. S. M. Condren

38. 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.
Dr. S. M. Condren

39. Geometries for Four Electron Pairs
Dr. S. M. Condren

40. Structure Determination by VSEPR
Ammonia, NH3
1. Draw electron dot structure
2. Count BP’s and LP’s = 4 H •• N
3. The 4 electron pairs are at the corners of a tetrahedron.
Dr. S. M. Condren

41. Structure Determination by VSEPR
Ammonia, NH3
The electron pair geometry is tetrahedral.
The MOLECULAR GEOMETRY, the
positions of the atoms, is PYRAMIDAL.
Dr. S. M. Condren

42. 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.
The molecular geometry is BENT.
Dr. S. M. Condren

43. Consequences of H2O Polarity
Dr. S. M. Condren

44. 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.
Dr. S. M. Condren

45. Molecular Geometries for Five Electron Pairs
All based on trigonal bipyramid
Molecular Geometries for Five Electron Pairs
Dr. S. M. Condren

46. Molecular Geometries for Six
All are based on the 8-sided octahedron
Molecular Geometries for Six
Electron Pairs
Dr. S. M. Condren

47. Bond Properties
What is the effect of bonding and structure on molecular properties?
Free rotation around C–C single bond
No rotation around C=C double bond
Dr. S. M. Condren

48. # of bonds between a pair of atoms
Bond Order
# of bonds between a pair of atoms
Double bond
Single bond
Acrylonitrile
Triple bond
Dr. S. M. Condren

49. Bond Order
Fractional bond orders occur in molecules with resonance structures.
Consider NO2- 1 1
The N—O bond order = 1.5
Dr. S. M. Condren

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

51. Dr. S. M. Condren

52. stronger, weaker, or the same strength?
Compare O-O and O=O. Is O=O expected to be
stronger, weaker, or the same strength?
Dr. S. M. Condren

53. longer, shorter, or the same length?
Is O=O expected to be
longer, shorter, or the same length?
Dr. S. M. Condren

54. 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
Net = ∆H = [S bondsbroken] – [S bondsformed]
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
∆Hfo (HCl(g)) = kJ/mol or -184 kJ
Dr. S. M. Condren

55. Water Boiling point = 100 ˚C Methane Boiling point = -161 ˚C
Molecular Polarity
Water Boiling point = 100 ˚C
Methane Boiling point = -161 ˚C
Why do water and methane differ so much in their boiling points?
Why do ionic compounds dissolve in water?
Dr. S. M. Condren

56. 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)
Dr. S. M. Condren

57. Bond Polarity Three molecules with polar, covalent bonds.
Each bond has one atom with a slight negative charge (-d) and another with a slight positive charge (+ d)
Dr. S. M. Condren

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

59. Electronegativity, 
 is a measure of the ability of an atom in a molecule to attract electrons to itself.
Dr. S. M. Condren

60. 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, .
Dr. S. M. Condren

61. Electronegativity Pauling Scale
relative attraction of an atom for electrons, its own and those of other atoms
same trends as ionization energy, increases from lower left corner to the upper right corner
fluorine: E.N. = 4.0
Dr. S. M. Condren

62. Electronegativities of the Elements
Dr. S. M. Condren

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

64. Molecular Polarity Molecules will be polar if a) bonds are polar AND
b) the molecule is NOT “symmetric”
All above are NOT polar
Dr. S. M. Condren

65. Compare CO2 and H2O. Which one is polar?
Polar or Nonpolar?
Compare CO2 and H2O. Which one is polar?
Dr. S. M. Condren

66. Covalent Bond Properties
electronegativity
nonpolar bonds => diff. EN = 0
polar bonds => diff. EN > 0
ionic bonds => diff. EN > 1.5
Dr. S. M. Condren

67. CH4 … CCl4 Polar or Not?
Only CH4 and CCl4 are NOT polar. These are the only two molecules that are “symmetrical.”
Dr. S. M. Condren