This document provides an overview of chemical bonding models including ionic bonding, covalent bonding, and metallic bonding. Ionic bonding involves the transfer of electrons between metals and nonmetals. Covalent bonding involves the sharing of electron pairs between two nonmetals. Metallic bonding occurs between metals and involves delocalized electrons within an electron sea. Lewis electron dot structures are used to represent electron arrangements around atoms and how they bond. Electronegativity differences between atoms determine bond polarity. Multiple bonding and resonance structures are also discussed.

1. CHEMICAL BONDING
Chapters 9.1-9.3, 9.5-9.6 &
10.1
• IONIC BONDS
• COVALENT BONDS
• LEWIS ELECTRON DOT
STRUCTURES
9-1

2. Chapter 9
Models of Chemical Bonding
9-2

3. Models of Chemical Bonding
9.1 Atomic Properties and Chemical Bonds
9.2 The Ionic Bonding Model
9.3 The Covalent Bonding Model
9.5 Between the Extremes:
Electronegativity and Bond Polarity
9.6 An Introduction to Metallic Bonding
9-3

4. Goals & Objectives
• See the following Learning
Objectives on pages 356 and 388.
• Understand these Concepts:
• 9.1-3, 6-10, 13-15; 10.1-4.
• Master these Skills: 9.1-2,6;
• 10.1-5.
9-4

5. Figure 9.1
9-5
A comparison of metals and nonmetals.

6. Types of Chemical Bonding
Ionic bonding involves the transfer of electrons and is
usually observed when a metal bonds to a nonmetal.
Covalent bonding involves the sharing of electrons and is
usually observed when a nonmetal bonds to a nonmetal.
Metallic bonding involves electron pooling and occurs
when a metal bonds to another metal.
9-6

7. Figure 9.2
9-7
Three models of chemical bonding.

8. CHEMICAL BONDS
• IONIC BONDING
– process involving the complete transfer of electrons
from one atom to another
– generally between metals and nonmetals
– for example Li metal and F a nonmetal
•
•
•
•
•
9-8
Li
+
F --> Li+ + F1s22s1
1s22s22p5
[He]
[Ne]
Li+ is isoelectronic with the He atom
F- is isoelectronic with the Ne atom

9. Lewis Electron-Dot Symbols
To draw the Lewis symbol for any main-group element:
•
Note the group number, which gives the number of
valence electrons.
•
Place one dot at a time on each of the four sides of the
element symbol.
•
Keep adding dots, pairing them, until all are used up.
Example:
Nitrogen, N, is in Group 5A and therefore has 5 valence electrons.
•
or • N
•
or
•
N•
•
••
9-9
or
•
• N•
••
••
••
•N•
•

10. Lewis Symbols and Bonding
For a metal, the total number of dots in the Lewis symbol
is the number of electrons the atom loses to form a cation.
For a nonmetal, the number of unpaired dots equals
- the number of electrons the atom gains to form an anion
- or the number it shares to form covalent bonds.
The octet rule states that when atoms bond, they lose,
gain, or share electrons to attain a filled outer level of 8
electrons (or 2, for H and Li).
9-10

11. Figure 9.4
Lewis electron-dot symbols for elements in Periods 2 and 3.
9-11

12. Lewis Electron Dot
Structures
• includes only the valence shell or
outermost shell electrons
• Draw the Lewis Electron Dot
Structures for the following atoms:
– Na
– Mg
–S
– Cl
9-12

13. 9-13

14. The Ionic Bonding Model
An ionic bond is formed when a metal transfers electrons
to a nonmetal to form ions, which attract each other to
give a solid compound.
The total number of electrons lost by the metal atom(s)
equals the total number of electrons gained by the
nonmetal atoms.
9-14

15. Figure 9.5 Three ways to depict electron transfer in the formation
of Li+ and F-.
Electron configurations
Li 1s22s1 + F 1s22p5
→ Li+ 1s2 + F-
1s22s22p6
Orbital diagrams
Li
↑↓
↑
+
1s
2s
2p
F
↑↓
↑↓
↑↓ ↑↓ ↑
1s
2s
2p
Li+ ↑↓
1s
2p
↑↓
↑↓
↑↓ ↑↓ ↑↓
1s
F-
2s
2s
2p
Lewis electron-dot symbols
Li+
+
•• F
••
••
••
9-15
••
Li•
••
•F
••

16. Sample Problem 9.1
Depicting Ion Formation
PROBLEM: Use partial orbital diagrams and Lewis symbols to depict
the formation of Na+ and O2− ions from the atoms, and
determine the formula of the compound formed.
PLAN: Draw orbital diagrams and Lewis symbols for Na and O
atoms. To attain filled outer levels, Na loses one electron and
O gains two. Two Na atoms are needed for each O atom so
that the number of electrons lost equals the number of
electrons gained.
SOLUTION:
••
••
•O
•
Na•
9-16
2Na+
•• 2+ O
••
••
••
Na•

17. Sample Problem 9.1
Na
↑
3s
Na
+
3p
O
↑
3s
↑↓
2s
3p
↑↓
↑↓ ↑↓ ↑↓
2s
2Na+ + O2-
2p
The formula is Na2O
9-17
↑↓ ↑
2p
↑

18. Figure 9.6
Na(s)
9-18
The exothermic formation of sodium bromide.
Br2(l)
NaBr(l)

19. Figure 9.7
9-19
The Born-Haber cycle for lithium fluoride.

20. Properties of Ionic Compounds
• Ionic compounds tend to be hard, rigid, and brittle, with
high melting points.
• Ionic compounds do not conduct electricity in the solid
state.
– In the solid state, the ions are fixed in place in the lattice and do
not move.
• Ionic compounds conduct electricity when melted or
dissolved.
– In the liquid state or in solution, the ions are free to move and
carry a current.
9-20

21. Figure 9.10
Solid ionic
compound
9-21
Electrical conductance and ion mobility.
Molten ionic
compound
Ionic compound
dissolved in water

22. Table 9.1 Melting and Boiling Points of Some Ionic Compounds
Compound
mp (°C)
bp (°C)
CsBr
636
1300
NaI
661
1304
MgCl2
714
1412
KBr
734
1435
CaCl2
782
>1600
NaCl
801
1413
LiF
845
1676
KF
858
1505
2852
3600
MgO
9-22

23. Lewis Electron Dot
Structures
• Draw the Lewis electron dot
structures for the reaction between Li
and F
• Draw the Lewis electron dot
structures for the reaction between
Mg and Cl
9-23
• Draw the Lewis electron dot
structures for the reaction between Li

24. 9-24

25. 9-25

26. CHEMICAL BONDS
• COVALENT BONDING
– involves the sharing of one or more
pairs of electrons between atoms
– usually occurs when two nonmetals
bond to each other
– Represent the formation of an H2
molecule from two H atoms by Lewis
electron dot formulas
9-26

27. Figure 9.12
9-27
Covalent bond formation in H2.

28. COVALENT BONDS
• Other simple compounds with
covalent bonds
– Halogens-F2, Cl2, Br2, I2
– Halogen acids-HF, HCl, HBr, HI
9-28

29. Figure 9.13
Distribution of electron density in H2.
At some distance (bond length),
attractions balance repulsions.
9-29
Electron density is high around
and between the nuclei.

30. Bonding Pairs and Lone Pairs
Atoms share electrons to achieve a full outer level of
electrons. The shared electrons are called a shared pair
or bonding pair.
The shared pair is represented as a pair of dots or a line:
••
H H
or
H–H
An outer-level electron pair that is not involved in
bonding is called a lone pair, or unshared pair.
••
9-30
or
•• ••
F–F
•• ••
••
••
••
••
•• ••
F F
•• ••

31. Properties of a Covalent Bond
The bond order is the number of electron pairs being
shared by a given pair of atoms.
A single bond consists of one bonding pair and has a bond order of 1.
The bond energy (BE) is the energy needed to
overcome the attraction between the nuclei and the
shared electrons. The stronger the bond the higher the
bond energy.
The bond length is the distance between the nuclei of
the bonded atoms.
9-31

32. Electronegativity and Bond Polarity
A covalent bond in which the shared electron pair is not
shared equally, but remains closer to one atom than the
other, is a polar covalent bond.
The ability of an atom in a covalent bond to attract the
shared electron pair is called its electronegativity.
Unequal sharing of electrons causes the more
electronegative atom of the bond to be partially negative
and the less electronegative atom to be partially positive.
9-32

33. POLAR AND NONPOLAR
COVALENT BONDS
• NONPOLAR COVALENT BONDS
– involve the equal sharing of the electron
pair between the two atoms
– H:H
nonpolar bond
• POLAR COVALENT BONDS
– involve the unequal sharing of the
electron pair between the two atoms
– H:Cl polar bond
9-33

34. Figure 9.20
Bonding between the models.
Polar covalent bonds are much
more common than either pure
ionic or pure covalent bonds.
9-34

35. Figure
9.22
Boundary ranges for classifying ionic character
of chemical bonds.
3.0
E
N
2.0
0.
0
9-35

36. Figure 9.21
9-36
The Pauling electronegativity (EN) scale.

37. Trends in Electronegativity
The most electronegative element is fluorine.
In general electronegativity decreases down a group as
atomic size increases.
In general electronegativity increases across a period
as atomic size decreases.
Nonmetals are more electronegative than metals.
9-37

38. Figure 9.22
9-38
Electronegativity and atomic size.

39. Depicting Polar Bonds
The unequal sharing of electrons can be depicted by a
polar arrow. The head of the arrow points to the more
electronegative element.
A polar bond can also be marked using δ+ and δ- symbols.
9-39

40. Figure 9.23
Electron density distributions in H2, F2, and HF.
In HF, the electron density shifts from H to F.
The H–F bond has partial ionic character.
9-40

41. Sample Problem 9.4
Determining Bond Polarity from EN Values
PROBLEM: (a) Use a polar arrow to indicate the polarity of each
bond: N–H, F–N, I–Cl.
(b) Rank the following bonds in order of increasing
polarity: H–N, H–O, H–C.
PLAN: (a) We use Figure 9.21 to find the EN values for each
element. The polar arrow points toward the more
electronegative element.
(b) The greater the EN between the atoms, the more polar
the bond.
SOLUTION: (a) The EN values are:
N = 3.0, H = 2.1; F = 4.0; I = 2.5, Cl = 3.0
N–H
9-41
F–N
I–Cl

42. Sample Problem 9.4
(b) The EN values are:
N = 3.0, H = 2.1; O = 3.5; C = 2.5
EN for H–N = 3.0 – 2.1 = 0.9
EN for H–O = 3.5 – 2.1 = 1.4
EN for H–C = 2.5 – 2.1 = 0.4
H-C < H-N < H-O
9-42

43. Predicting Polarity
• Predict the polarity of each of the
following compounds:
• LiCl
• CO
• N2
9-43

44. 9-44

45. Metallic Bonding
The electron sea model of metallic bonding proposes that:
• All metal atoms in the sample contribute their valence
electrons to form a delocalized electron “sea”.
• The metal “ions” (nuclei with core electrons) lie in an
orderly array within this mobile sea.
• All the atoms in the sample share the electrons.
• The metal is held together by the attraction between the
metal “cations” and the “sea” of valence electrons.
9-45

46. Properties of Metals
• Metals are generally solids with moderate to high melting
points and much higher boiling points.
– Melting points decrease down a group and increase across a
period.
• Metals can be shaped without breaking.
– The electron sea allows the metal ions to slide past each other.
• Metals are good conductors of electricity in both the solid
and liquid states.
– The electron sea is mobile in both phases.
• Metals are good conductors of heat.
9-46

47. Table 9.5 Melting and Boiling Points of Some Metals
Element
bp ( C)
Lithium (Li)
180
1347
Tin (Sn)
232
2623
Aluminum (Al)
660
2467
Barium (Ba)
727
1850
Silver (Ag)
961
2155
Copper (Cu)
1083
2570
Uranium (U)
9-47
mp ( C)
1130
3930

48. Figure 9.28
Melting points of the Group 1A(1) and Group 2A(2) metals.
9-48

49. Figure 9.29
9-49
Why metals dent and bend rather than crack.

50. WRITING LEWIS
ELECTRON DOT
STRUCTURES
• N - A = S Rule
• N = total number of electrons needed
to achieve a rare gas configuration
• A = total number of electrons
available
• S = total number of electrons shared
– must be an even number
– must be less than or equal to the
number of electrons available
9-50

51. WRITING LEWIS
ELECTRON DOT
STRUCTURES
9-51
• The central atom in a molecule or
polyatomic ion is the atom requiring
the greatest number of additional
electrons to achieve a rare gas
configuration.
• When deciding between two
elements in the same group, the less
electronegative atom is usually
chosen.
• H is never the central atom.

52. WRITING LEWIS
ELECTRON DOT
STRUCTURES
• Write Lewis electron dot structures
and dash formulas for water, H2O.
• Write Lewis electron dot structures
and dash formulas for ammonia, NH3.
• Write Lewis electron dot structures
and dash formulas for methane, CH4.
9-52

53. 9-53

54. 9-54

55. 9-55

56. Sample Problem 10.3
Writing Lewis Structures for Molecules
with Multiple Bonds
PROBLEM: Write Lewis structures for the following:
(a) Ethylene (C2H4), the most important reactant in the
manufacture of polymers
(b) Nitrogen (N2), the most abundant atmospheric gas
PLAN: After following steps 1 to 4 we see that the central atom does not
have a full octet. We must therefore add step 5, which involves
changing a lone pair to a bonding pair.
SOLUTION:
(a) C2H4 has 2(4) + 4(1) = 12 valence e-. H can have only one bond
per atom.
9-56

57. Sample Problem 10.3
(b) N2 has 2(5) = 10 valence e-.
9-57

58. WRITING LEWIS
ELECTRON DOT
STRUCTURES
• Write a Lewis electron dot structure
for oxygen, O2.
• Write a Lewis electron dot structure
for nitrogen, N2.
• Write a Lewis electron dot structure
for carbon dioxide, CO2.
9-58

59. 9-59

60. 9-60

61. 9-61

62. Resonance Structures
O3 can be drawn in 2 ways:
These are two different reasonable Lewis structures for
the same molecule.
Neither structure depicts O3 accurately, because in reality
the O-O bonds are identical in length and energy.
9-62

63. Resonance Structures
The structure of O3 is shown more correctly using both
Lewis structures, called resonance structures.
A two-headed resonance arrow is placed between them.
Resonance structures have the same relative placement
of atoms but different locations of bonding and lone
electron pairs.
9-63

64. The Resonance Hybrid
A species like O3, which can be depicted by more than
one valid Lewis structure, is called a resonance hybrid.
Resonance forms are not real bonding depictions.
O3 does not change back and forth between its two
resonance forms.
The real structure of a resonance hybrid is an average
of its contributing resonance forms.
9-64

65. Sample Problem 10.4
Writing Resonance Structures
PROBLEM: Write resonance structures for the nitrate ion, NO3− and
find the bond order.
PLAN: Write the Lewis structure, remembering to add 1e- to the total
number of valence e- for the -1 charge. We apply Step 5 to
form multiple bonds. Since multiple bonds can be formed in
more than one location, there are resonance forms.
SOLUTION:
Nitrate has [1 x N(5e-)] + [3 x O(6e-)] + 1e-] = 24 valence eAfter Steps 1-4:
9-65

66. Sample Problem 10.4
Step 5. Since N does not have a full octet, we change a lone pair from O
to a bonding pair to form a double bond.
4 shared electron pairs
Bond order =
= 1⅓
3 bonded-atom pairs
9-66

67. Formal Charge
Formal charge is the charge an atom would have if all
electrons were shared equally.
Formal charge of atom =
# of valence e- - (# of unshared valence e- + ½ # of shared valence e-)
For OA in resonance form I, the formal charge is given by
6 valence e- - (4 unshared e- + ½(4 shared e-) = 6 – 4 – 2 = 0
9-67

68. Formal Charge
Formal charges must sum to the actual charge on the
species for all resonance forms.
OA [6 – 4 – ½(4)] = 0
OB [6 – 2 – ½(6)] = +1
OC [6 – 6 – ½(2)] = -1
OA [6 – 6 – ½(2)] = -1
OB [6 – 2 – ½(6)] = +1
OC [6 – 4 – ½(4)] = 0
For both these resonance forms the formal charges sum to zero,
since O3 is a neutral molecule.
9-68

69. Choosing the More Important Resonance Form
• Smaller formal charges (positive or negative) are
preferable to larger ones.
• The same nonzero formal charges on adjacent atoms are
not preferred.
Avoid like charges on adjacent atoms.
• A more negative formal charge should reside on a more
electronegative atom.
9-69

70. Example: NCO− has 3 possible resonance forms:
+2 0
-1
-1
0
0
0
0
-1
Resonance forms with smaller formal charges are preferred.
Resonance form I is therefore not an important contributor.
A negative formal charge should be placed on a more electronegative
atoms, so resonance form III is preferred to resonance form II.
The overall structure of the NCO- ion is still an average of all three
forms, but resonance form III contributes most to the average.
9-70

71. WRITING LEWIS
ELECTRON DOT
STRUCTURES
• RESONANCE STRUCTURES are
used to represent molecules when
two or more dot representations are
needed to describe the molecule.
• Write Lewis electron dot structures
for the resonance forms expected for
sulfur trioxide, SO3.
• Write dash formulas for sulfur
trioxide.
9-71

72. 9-72

73. 9-73

74. Exceptions to the Octet Rule
Molecules with Electron-Deficient Atoms
B and Be are commonly
electron-deficient.
Odd-Electron Species
A molecule with an odd number of electrons is
called a free radical.
9-74

75. Exceptions to the Octet Rule
Expanded Valence Shells
An expanded valence shell is only possible for nonmetals
from Period 3 or higher because these elements have
available d orbitals.
9-75

76. MOLECULES THAT DO
NOT OBEY THE N-A=S
RULE
• COVALENT COMPOUNDS OF Be
• COVALENT COMPOUNDS OF B
• Draw the Lewis electron dot structure
for the BBr3 molecule.
• Draw the Lewis electron dot structure
for the BeCl2 molecule.
9-76

77. 9-77

78. 9-78