1. Chapter 8:
Chemical Bonding

2. Questions for Consideration
How can we classify the types of bonding of different compounds?
What is the nature of the bonding in ionic compounds?
What is the nature of the bonding in molecular compounds?
How do we predict and explain the shapes of molecular compounds and polyatomic ions?

3. Chapter 8 Topics: Types of Bonds Ionic Bonding Covalent Bonding
Bonding in Carbon Compounds
Shapes of Molecules

4. Introduction
How and why do atoms come together (bond) to form compounds?
Why do different compounds have such different properties?
What do molecules look like in three dimensions?

5. 8.1 Types of Bonds Chemical bond Two types of chemical bonds
A force that holds atoms together in a molecule or compound
Two types of chemical bonds
Ionic Bonds
Covalent Bonds
Figure 8.2
Figure 8.2

6. Table 8.1 Properties of Two Carbon Compounds
Figure 8.2

7. General Properties of Ionic and Covalent Substances

8. Ionic and Covalent
In ionic compounds, ions are held together by electrostatic forces – forces between oppositely charged ions.
In molecular compounds, atoms are held together by covalent bonds in which electrons are shared.
Figure 8.2

9. Ionic Bonds
A bond created by electrostatic attraction between oppositely charged ions
Occurs between a metal and a nonmetal
Electrons transferred between the cation (positively charged ion) and the anion (negatively charged ion)
Extremely strong bonds
Figure 8.7

10. Covalent Bonds
A bond created by the sharing of electrons between atoms
Occurs between two nonmetals (resulting in a neutral overall charge)
Electrons not transferred in this case
Electrons typically shared in pairs
Weaker bonds than ionic bonds
Figure 8.2

11. Activity: Identifying Types of Bonding
Identify the type of bonding in each of the following substances:
NaF
ClO2
FeSO4
SO2
Ca(ClO2)2

12. Activity Solutions: Identifying Types of Bonding
Identify the type of bonding in each of the following substances:
NaF – Ionic bonding (metal + nonmetal)
ClO2 – Covalent bonding (2 nonmetals)
FeSO4 – Ionic bonding between the metal and polyatomic ion; covalent bonding between the nonmetals in the polyatomic ion
SO2 – Covalent bonding
Ca(ClO2)2 - Ionic bonding between the metal and polyatomic ion; covalent bonding between the nonmetals in the polyatomic ion

13. Polar vs. Nonpolar Two general types of covalent bonds: Polar covalent
Unequal sharing (or a partial transfer) of electrons
Occurs when different elements are covalently bonded to one another
Why different elements?
Because different elements have different electronegativities
Nonpolar covalent
Equal sharing (no transfer) of electrons
Occurs only when all of the atoms in a molecule belong to the same element

14. Polar vs. Nonpolar Polar covalent bonds are:
Typically shorter bonds
Stronger bonds due to their increased ionic character
Nonpolar covalent bonds are:
Typically longer bonds
Weaker bonds
Polarity
Occurs in polar covalent molecules
Polarity is the degree of transfer of electrons in a covalently bonded molecule composed of different element’s atoms.

15. Classification of Bonding
Figure 8.3

16. Polar Covalent Bonds
In polar covalent bonds, electrons are not shared equally. This result in unequal sharing which can be described as partial electron transfer.
In a polar bond there are partial charges on the atoms sharing electrons.
Figure 8.4

17. Electronegativity Ability of an atom to attract bonding electrons
Proposed by Linus Pauling in the early 1930’s
A difference in electronegativity between the atoms in a covalent bond results in:
A polar covalent bond
Increased ionic character
The greater the difference in electronegativity, the greater the ionic character and the more polar the bond that joins the atoms.
Decreased bond length and increased bond strength
No difference in electronegativity between atoms in a covalent bond results in a nonpolar covalent bond.

18. Electronegativity
The electronegativity scale tells us which elements have a greater pull on electrons in a covalent bond.
The more electronegative element pulls bonding electrons more strongly and obtains a partial negative charge ().
The less electronegative element loses electron density in the bond and therefore obtains a partial positive charge (+).
Figure 8.4

19. Electronegativity Values
Figure 8.5

20. Electronegativity Trends
Figure 8.6

21. Electronegativity Trends
The difference in electronegativity between metals and nonmetals is so large, that the electrons are transferred, not shared.
The greater the electronegativity difference, the more polar the bond.
Si-F > N-F> O-F >F-F
What partial charges go
on each atom?
Figure 8.6

22. Activity: Polar Bonds
Which of the following molecules have polar bonds? If a bond is polar, which atom has a partial negative charge?
SO2 N2
H2S
CCl4 O3

23. Activity Solutions: Polar Bonds
Which of the following molecules have polar bonds? If a bond is polar, which atom has partial negative charge?
SO2 – Polar covalent bonds; O is more electronegative and has a partial negative charge
N2 – Nonpolar covalent bonds
H2S – Polar covalent bonds; S is more electronegative and has a partial negative charge
CCl4 – Polar covalent bonds; Cl is more electronegative and has a partial negative charge
O3 – Nonpolar covalent bonds

24. 8.2 Ionic Bonding
Each element immediately following a noble gas is a metal.
Metals lose electrons, forming a positive charge, to become cations.
Each element immediately preceding a noble gas is a nonmetal.
Nonmetals gain electrons, forming a negative charge, to become anions.
Formation of ions and ionic bonds relates to an element’s electron configuration.
Many main-group elements either lose or gain electrons to become isoelectronic with a noble gas (i.e. have the same electron configuration). As ions, they are known as the common ions.

25. Formation of Ionic Bonds
To understand ionic bonding, we can think about how ions and ionic bonds would be formed from neutral atoms of the elements.
Consider the formation of NaCl from its elements:
Electrons are transferred from the metal to the nonmetal to form ions, each with a noble-gas electron configuration.
Figure 6.6

26. Lewis Dot Symbols Lewis Dot symbol Electron dot symbol
Dots placed around an element’s symbol represent valence electrons
Pair electrons as needed
Octet rule
Tendency of an atom to achieve an electron configuration having 8 valence electrons
Same as the electron configuration of a noble gas
The 8 electrons exist in 4 pairs
Ions achieve 8 electrons by losing or gaining electrons

27. Lewis Symbols
Lewis Symbols help us to focus on the valence electrons – those that can participate in bonding.

28. Activity: Lewis Symbols for Ions
Write the Lewis symbols for beryllium and nitrogen ions. Then write a formula for the compound that would form between them, using their Lewis symbols.

29. Activity Solutions: Lewis Symbols for Ions
Write the Lewis symbols for the beryllium and nitrogen ions. Then write a formula for the compound that would form between them, using their Lewis symbols.
. .
symbol

30. Ionic Compounds
Ions of like charge repel each other in ionic compounds, and opposite charged ions attract. This results in a 3-dimensional regular pattern called a crystal lattice.
Figure 8.7

31. Structures of Ionic Compounds
Crystal lattice
The pattern obtained when an ion, represented as a charged sphere, exerts a force equally in all directions.
Thus, ions of equal and opposite charge surround it.
Cations and anions must come into contact for a crystal lattice to form.
Figure 8.10
fluorite, CaF2

32. Structures of Ionic Compounds
Sodium Chloride, NaCl, Crystal
Figure 8.8

33. Cesium Chloride Crystal
Figure 8.9

34. Calcium Fluoride Crystal
Figure 8.10

35. Properties of Ionic Solids
Why are crystalline ionic solids hard and brittle?
Why are they poor conductors of electricity in the solid state?
Figure 8.11

36. Structures of Ionic Crystals
Ions are arranged in a regular geometric pattern that maximizes the attractive forces and minimizes the repulsive forces.
Hard and brittle
Can shatter if struck forcefully
The charges and sizes of ions largely determine the characteristic patterns of ionic crystals
Figure 8.11

37. 8.3 Covalent Bonding
When two nonmetals form a bond, the bond is covalent. They are both close to the noble-gas electron configuration, so sharing will allow both to obtain it.
In a covalent bond, each shared electron interacts simultaneously with two nuclei.
Figure 8.12

38. Carbon Dioxide – Covalent bonds
The atoms of CO2 molecules are held together by strong covalent bonds. No bonds connect the molecules, so CO2 molecules separate from each other into the gas state at room temperature.
Figure from p. 27
Figure 8.13

39. The Octet Rule
Just as in ionic bonding, covalent bonds are formed so that each atom can have the noble-gas electron configuration. Noble gases have 8 valence electrons, an octet.
Figure 8.14

40. The Octet Rule Octet rule
Tendency of an atom to achieve an electron configuration having 8 valence electrons
Same as the electron configuration of a noble gas
Covalently bonded atoms achieve 8 valence electrons by sharing electrons
The 8 electrons exist in 4 pairs
H atoms bond with other atoms to obtain a total of 2 electrons like He.

41. Lewis Formulas for the Diatomic Elements
How does hydrogen obtain a noble-gas electron configuration?
Figure 8.15

42. The Halogens Do the atoms in each of these molecules have an octet?
Why do the halogens exist as diatomic molecules?
Figure 8.16

43. Multiple Bonds How many valence electrons does an oxygen atom have?
How many does it need to obtain an octet?
Can a single covalent bond allow each oxygen atom to satisfy the octet rule?
O2 has a double bond, two pairs of shared electrons
Figure from p. 315
Figure 8.17

44. Multiple Bonds How many valence electrons does a nitrogen atom have?
How many does it need to obtain an octet?
N2 has a triple bond, three pairs of shared electrons
Figure 8.17

45. P4 and S8 How do phosphorus and sulfur obtain an octet in P4 and S8?
Figure 8.18

46. Activity: Valence Electrons and Number of Bonds
How many bonds do each of the following atoms tend to form? H Cl O N C 1 2 3 4

47. Activity: Lewis Structures
Draw the Lewis structures for each of the following based on the number of bonds that each tends to form. Remember to include the nonbonding electrons so that all have octets (except H).
C2H6
C2H4
C2H2
HCN

48. Activity Solutions: Lewis Structures
Draw the Lewis structures for each of the following based on the number of bonds that each tends to form.
Carbon forms 4 bonds and hydrogen forms 1:
Hydrogen forms 1 bond, carbon forms 4 bonds, and nitrogen forms 3 bonds:

49. Extra Activity: Valence Electrons and Number of Bonds
Draw a Lewis structure for each of the following based on how many bonds each tends to form. Remember to include the nonbonding electrons so that all have octets (except H).
H2CO
NF3
H2O

50. Steps for Writing Lewis Structures
Write an atomic skeleton.
Sum the valence electrons from each atom to get the total number of valence electrons.
Place two electrons, a single bond, between each pair of bonded atoms.
Add remaining electrons to complete the octet of each outer atom and then to the central atom if necessary and if there are electrons available.
If necessary to satisfy the octet rule, shift unshared electrons from non-bonded positions on atoms with completed octets to positions between atoms to make double or triple bonds.

51. Activity: Writing Lewis Structures
Write a Lewis formula for the formaldehyde molecule, CH2O.

52. Activity Solution: Writing Lewis Structures
Write an atomic skeleton for CH2O:
Sum the valence electrons from each atom to get the total number of valence electrons.
Carbon is in Group IVA (14), so it has 4 valence electrons. Each hydrogen contributes 1 valence electron (H is in Group IA (1)). Oxygen contributes 6 valence electrons because it is in Group VIA (16).
Total number of valence electrons = 4 + (1 × 2) + 6
= 12

53. Activity Solution: Writing Lewis Structures
Place a single bond between each pair of bonded atoms.
Add the remaining electrons to complete the octet of each outer atom, then to the central atom if any are left (not the case here).
All the electrons are used, but carbon does not have an octet, so shift electrons to make a double bond.

54. Activity: Writing Lewis Structures
Write a Lewis formula for hydrogen cyanide, HCN.

55. Activity Solution: Writing Lewis Structures
Write an atomic skeleton for HCN:
Sum the valence electrons from each atom to get the total number of valence electrons.
Carbon is in Group IVA (14), so it has 4 valence electrons. Hydrogen contributes 1 valence electron (H is in Group IA (1)). Nitrogen contributes 5 valence electrons because it is in Group VA (15).
Total number of valence electrons = = 10

56. Activity Solutions: Writing Lewis Structures
Place a single bond between each pair of bonded electrons.
Add the remaining electrons to complete the octet of the outer atoms.
Since carbon does not have an octet, shift electrons to make a double bond, or a triple bond, until carbon has an octet.

57. Activity: Writing Lewis Structures
Write Lewis structures for each of the following:
CO2
SO2
NO2
HNO3

58. Resonance
Sometimes a single Lewis structure does not give an accurate picture of the bonding.
For example, there are two Lewis structures that represent the ozone, O3 , molecule.
The actual molecule is a blend of these two:
Both O−O bonds are the same length with partial double bond character
Figure from p. 313

59. Resonance
Like a single Lewis structure, each alone shows some aspect of the shell.
Like a set of resonance structures, the pair of photos together gives a better representation of the shell.
Figure 8.19

60. Activity: Resonance Structures
Write resonance structures for N2O.
The atom arrangement is N N O.

61. Activity Solution: Resonance Structures
Write resonance structures for N2O.
Write an atomic skeleton for N2O:
N N O
Sum the valence electrons from each atom to get the total number of valence electrons.
Nitrogen is in Group VA (15), so it has 5 valence electrons. Oxygen contributes 6 valence electrons because it is in Group VIA (16).
Total number of valence electrons = (5 × 2) + 6 = 16

62. Activity Solution: Resonance Structures
Next, form a bond between each pair of atoms.
Then add electrons to give the outer atoms an octet; this uses all 16 electrons.
Since the central nitrogen doesn’t have an octet, shift electrons to make a double bond.
The central nitrogen still doesn’t have an octet, so shift another pair of electrons: or

63. Activity Solution: Resonance Structures
The two resonance structures are shown connected by a double-headed arrow:

64. Activity: Resonance Structures
Which of the following have resonance?
For those that do, how many resonance structures do they have?
NO3
HNO3
CH4
H2CO
Yes, 3
Yes, 2 No

65. Exceptions to the Octet Rule
Molecules with an odd number of electrons
e.g. NO
Incomplete octets
e.g. BH3 or BF3
Expanded octets
e.g. SF4 or SF6

66. 8.4 Bonding in Carbon Compounds
Figure 8.21
Carbon has:
Four valence electrons
The ability to form four bonds
The ability to bond to itself
Very strong bonds when bonded to itself
Carbon molecules are ubiquitous in nature.

67. Hydrocarbons
Figure 8.20

68. Hydrocarbons Compounds containing hydrogen and carbon
Aliphatic hydrocarbons
A class in which the bonds are all localized single, double, and triple bonds
Alkanes
Hydrocarbons which contain only carbon-carbon single bonds
Alkenes
Hydrocarbons which contain at least one carbon-carbon double bond
Alkynes
Hydrocarbons which contain at least one carbon-carbon triple bond

69. Hydrocarbons Aromatic hydrocarbons
A class of hydrocarbons which has carbon atoms arranged in a six-atom ring with alternating single and double bonds
Delocalized structures
Figure 8.22

70. Functional Groups Functional group
A group that is introduced into or substituted in a hydrocarbon chain
Gives the hydrocarbon its characteristic properties
The group has a heteroatom, an atom other than C and H
Typically O, S, and N
Alcohol
A hydroxyl group (-OH) replaces a hydrogen atom in the formula for a hydrocarbon
Figure 8.23

71. Functional Groups in Hydrocarbons

72. Activity: Functional Groups
Identify the functional group in each of the following molecules:
CH3OH
(CH3)2CO
CH3CH2CO2H
CH3CH2NH2
CH3CH2CHO
CH3CO2CH2CH
alcohol
ketone
carboxylic acid
amine
aldehyde
ester

73. 8.5 Shapes of Molecules
The shape of a molecule influences many of its properties including taste and smell.
Sweet substances are often of a shape similar to that of glucose. It’s –H and –OH groups fit into a taste receptor site on the tongue.
Figure 8.23

74. Odor Receptors
Figure 8.24

75. Predicting Shapes of Molecules
A simple model used with Lewis structures allows us to predict shapes of molecules.
Valence-Shell Electron-Pair Repulsion Theory (VSEPR theory) is based on the fact that negative charges repel one another.
Valence electron pairs (bonding or nonbonding) repel one another and take up positions that maximize their distance and angles between them.

76. VSEPR (CH4)
The methane molecule (CH4) has four pairs of electrons around the central atom (all bonding):
Maximizing their distance gives bond angles of 109.5°. H 
H―C―H
Figure from p. 328

77. Parent Structures
Figure 8.25

78. Parent Structures
When determining the parent structure and bond angles for a molecule or polyatomic ion, each of the following is counted as an electron domain:
Nonbonding electron pair
Single bond
Double bond
Triple bond
Multiple bonds don’t repel and separate
from each other so they are counted only
once even though they contain more than one pair.

80. Activity: Parent Structures and Bond Angles
Determine the parent structure and bond angles for each of the following.
NH3
CO2
SO2
BH3
H2O
Tetrahedral, 109.5o
Linear, 180o
Trigonal planar, 120o
Figure 8.27

81. Shapes of Molecules and Ions
The actual shape of a molecule includes only the atoms.
To envision the shape from a parent structure, make the nonbonding electrons invisible (but they are still there).
What is the shape of a water molecule?
Figure 8.27

82. Shapes of Molecules and Ions
Why do CO2 and SO2 have different shapes?
Figure 8.28

83. Steps for VSEPR Structures
Draw a Lewis formula.
Count the number of atoms bonded to the central atom, and count unshared pairs on the central atom.
Add the number of atoms and the number of unshared electron pairs around the central atom. The total indicates the number of electron domains and thus the parent structure.
The molecular shape is derived from the parent shape by considering only the positions in the structure occupied by bonded atoms.

86. Molecular Shapes
Trigonal planar and trigonal pyramidal molecular shapes are different from one another.
Figure 8.26

87. Activity: VSEPR Structures
What is the shape of the nitrite ion (NO2−)? What is the O − N − O bond angle?

88. Activity Solution: VSEPR Structures
Draw a Lewis formula.
Count the number of atoms bonded to the central atom and count unshared pairs on the central atom.
N, which is the central atom in this case, has 2 atoms bonded to it, and 1 unshared pair.

89. Activity Solutions: VSEPR Structures
If we add the number of atoms and unshared pairs around the central atom, we get the number of electron domains. There are 3 so this indicates that the parent structure is trigonal planar. o

90. Activity Solutions: VSEPR Structures
The molecular shape is determined by considering only the positions of atoms. This gives a bent shape with a bond angle of about 120°.
120°

91. Describing Shapes of Molecules with More than One Central Atom
Glycine: H2NCH2CO2H
Molecular shapes are important in living systems.
Glycine molecules are typically found in proteins or gelatins.
Draw the Lewis structure.
Determine the parent structure, remembering to consider all the electron domains.
Determine the bond angles at each central atom (N, C, C, and O)

92. Describing Shapes of Molecules with More than One Central Atom
Glycine: H2NCH2CO2H
Draw the Lewis structure.
Determine the parent structure, remembering to consider all the electron domains.
Determine the bond angles at each central atom (N, C, C, and O).
Figure 8.29

93. Activity: What are the bond angles?
Figure 8.21
109.5o

94. Activity: What are the bond angles?
Figure 8.22
120o

95. Natural Applications of VSEPR Theory
Molecular shapes are important in living systems.
Glycine molecules are typically found in proteins or gelatins.
Figure 8.29

96. Natural Applications of VSEPR Theory
Molecular shapes are important in living systems.
Heme molecule
Oxygen is carried throughout the body via red blood cells containing heme molecules.
Histidine, an amino acid in the heme molecule, just fits into the space next to the oxygen molecule.
Figure 8.30

97. Polarity of Molecules Diatomic molecules Polyatomic molecules
Polarity lies along the plane of the bond
Polyatomic molecules
A nonpolar molecule is one that has all nonpolar bonds or one that has polar bonds that cancel out
Bonds that cancel out have equal polarities in opposite directions
This happens when:
A central atom has no unshared electrons
The atoms around the central atom all have the same electronegativity
A polar molecule is one that has polar bonds that DO NOT cancel out

98. Polarity of Molecules Which of these substances is a polar substance?
Figure 8.32

99. Molecular Polarity
Molecules with polar bonds can be polar under certain conditions.
A diatomic molecule with a polar bond is polar since there is only one bond.
A molecule or polyatomic ion with polar bonds will be
polar if its geometry causes an asymmetric distribution of partial charge. (e.g. H2O)
Nonpolar if its geometry is symmetrical enough to allow cancellation of the polar bonds. (e.g. CO2)
A polyatomic molecule with nonpolar bonds cannot be polar. O3 and NBr3 have no polar bonds.
Figure 8.4

100. Nonpolar Molecules with Polar Bonds
The molecules SO3 and CCl4 have polar bonds, but their symmetrical shapes with no nonbonding electrons on the central atom cause them to be nonpolar molecules.
Trigonal Planar
Figure 8.33
Tetrahedral

101. Asymmetric Molecules
In asymmetric molecules like these, the atoms surrounding the central atom are not identical and have different bond polarities. There is a net polarity to the molecules.
Figure 8.34

102. Activity: Which of the following are polar molecules?
Figure from p. 341

103. All Hydrocarbons are Nonpolar
Figure 8.21
Figure 8.22

104. Activity: Polarity of Molecules
Predict whether C2H6, NO2, CO2, SO2, and SO3 are polar or nonpolar molecules.

105. Activity Solutions: Polarity of Molecules
Predict whether C2H6, NO2, CO2, SO2, and SO3 are polar or nonpolar molecules.
C2H6 tetrahedral at each C; nonpolar
NO2 bent; polar
CO2 linear; nonpolar
SO2 bent; polar
SO3 trigonal planar; nonpolar

106. Extra Activity: Which of the following are polar molecules?
NH3
BH3
CS2
H2S O3
CH2Cl2
polar
nonpolar

107. Polarity and Properties
Molecular polarity is important in many properties of molecules. One example is solubility, which is discussed in more detail in Chapter 11.
A rule of thumb for solubility is “like dissolves like”, which refers to similarity in polarity.

108. “Like” Dissolves “Like”
Ionic salts and polar liquids dissolve better in polar liquids than in nonpolar liquids.
Nonpolar liquids dissolve better in other nonpolar liquids than in polar liquids.
Figure 8.35