1. Chemical Bonding I: The Covalent Bond
Chapter 9

2. Problems from pages
1-4, 5-26, 28, 30, 32, 34-36, 38, 40, 42, 44-46, 48, 50, 52, 54, 66

3. Test on Chapters 7 and 8 April 6
EH Assignment
Due April 4
Unit 11
Minimum score
Sec 1 Size of Atoms and Ions 90
Sec 2 Ionization Energy & EA
Sec 3 Metals, Nonmetals, Families
Sec 4 Valence Electrons & Charge
Sec 5 Acidity and Basicity 75
Test on Chapters 7 and 8 April 6

4. Valence electrons are the outer shell electrons of an
atom. The valence electrons are the electrons that
particpate in chemical bonding.
Group
# of valence e-
e- configuration 1A 1
ns1 2A 2
ns2 3A 3
ns2np1 4A 4
ns2np2 5A 5
ns2np3 6A 6
ns2np4 7A 7
ns2np5
9.1

5. 9.1

6. The Ionic Bond - - - Li+ F Li + F 1s22s1 1s22s22p5 [He] 1s2 1s22s22p6
[Ne] Li
Li+ + e-
e- + F - F -
Li+ +
Li+
9.2

7. Covalent Bonding I
Chemists use several different models to describe the covalent
bond. These range from very simple to quite complex.
The simplest is the Lewis Structure model which is discussed
in detail in this chapter. It is very useful and successfully
predicts the formulas and bonding in many molecules and
polyatomic ions.
Valence electrons are represented as dots. The octet rule is used.
An atom tends to form bonds until it is surrounded by the same
number of electrons as a noble gas.

8. Covalent Bonding II In the Lewis structure model valence electrons are
presented as dots. The covalent bond is represented by a
shared pair (bonding pair) of electrons. An electron pair
that is not involved in bonding is called a lone pair or
unshared pair or nonbonding pair.
Bond Order: single, double, triple bonds.
But there are many things this model cannot explain. In the
next chapter the Valence Bond model and Molecular Orbital
models are presented.

9. Why should two atoms share electrons?
A covalent bond is a chemical bond in which two or more electrons are shared by two atoms.
Why should two atoms share electrons?
7e-
7e-
8e-
8e- F F + F
Lewis structure of F2
lone pairs F
single covalent bond
single covalent bond F
9.4

10. Lewis structure of water
single covalent bonds
2e-
8e-
2e- H + O + H O H or
Double bond – two atoms share two pairs of electrons
8e-
8e-
8e-
double bonds O C or O C
double bonds
Triple bond – two atoms share three pairs of electrons
8e-
triple bond N
8e- or N
triple bond
9.4

11. Lengths of Covalent Bonds
Bond Type
Bond Length
(pm)
C-C
154
CC
133
CC
120
C-N
143
CN
138
CN
116
Bond Lengths
Triple bond < Double Bond < Single Bond
9.4

12. Last Test – April 26? EH Assignment Due April 15 Unit 12 Minimum score
Sec 1 Electronegativity and Bond Polarity 90
Sec 2 Lewis Dot Diagrams 88
Sec 3 Shapes of Ions and Molecules
later
Sec 4 Resonance and Formal Charge
Sec 5 Review of Molecular Shapes
Last Test – April 26?

13. Electronegativity and Bond Polarity
Some bond are purely ionic - a complete transfer of electrons has occurred. NaCl
Some bonds are pure covalent - the bonding electrons are
shared equally. H2 or Cl2
In many cases the electrons are shared, but not equally.
These bonds are called polar-covalent are considered to be
partially ionic and partially covalent. HCl
Electronegativity: the relative ability of a bonded atom to
attract the shared electrons.

14. Polar covalent bond or polar bond is a covalent bond with greater electron density around one of the two atoms H F
electron rich
region
electron poor
region
e- poor
e- rich F H d+ d-
9.5

15. Electron Affinity - measurable, Cl is highest
Electronegativity is the ability of an atom to attract toward itself the electrons in a chemical bond.
Electron Affinity - measurable, Cl is highest
X (g) + e X-(g)
Electronegativity - relative, F is highest
9.5

16. 9.5

17. Classification of bonds by difference in electronegativity
Bond Type
Covalent
 2
Ionic
0 < and <2
Polar Covalent
Increasing difference in electronegativity
Covalent
share e-
Polar Covalent
partial transfer of e-
Ionic
transfer e-
9.5

18. Classify the following bonds as ionic, polar covalent,
or covalent: The bond in CsCl; the bond in H2S; and
the NN bond in H2NNH2.
Cs – 0.7
Cl – 3.0
3.0 – 0.7 = 2.3
Ionic
H – 2.1
S – 2.5
2.5 – 2.1 = 0.4
Polar Covalent
N – 3.0
N – 3.0
3.0 – 3.0 = 0
Covalent
9.5

19. Writing Lewis Structures
Draw skeletal structure of compound showing what atoms are bonded to each other. Put least electronegative element in the center.
Count total number of valence e-. Add 1 for each negative charge. Subtract 1 for each positive charge.
Complete an octet for all atoms except H until all are complete or you run out of electrons.
Form double and triple bonds on central atom if needed to complete octet.
9.6

20. ( ) - - Two possible skeletal structures of formaldehyde (CH2O) H C O
An atom’s formal charge is the difference between the number of valence electrons in an isolated atom and the number of electrons assigned to that atom in a Lewis structure.
formal charge on an atom in a Lewis structure =
total number of valence electrons in the free atom -
total number of nonbonding electrons - 1 2
total number of bonding electrons ( )
The sum of the formal charges of the atoms in a molecule or ion must equal the charge on the molecule or ion.
9.7

21. Formal Charge
Formal charge: the charge an atom would have if the bonding
electrons were shared equally.
The formal charges must add up to equal the actual charge on
the species.
Atoms with a “normal” number of bonds have zero formal
charge.
Normal number of bonds - Cl, O, N, C, H

22. Formal Charge and Lewis Structures
For neutral molecules, a Lewis structure in which there are no formal charges is preferable to one in which formal charges are present.
Lewis structures with large formal charges are less plausible than those with small formal charges.
Among Lewis structures having similar distributions of formal charges, the most plausible structure is the one in which negative formal charges are placed on the more electronegative atoms.
Which is the most likely Lewis structure for CH2O? H C O -1 +1 H C O
9.7

23. For some molecules and ions no single Lewis
Resonance
For some molecules and ions no single Lewis
Structure can be drawn that agrees with experiment.
In cases like this we can:
Go on to a more sophisticated model of
the covalent bond or
2. We can try to patch up the simple Lewis
structure model to make it work.

24. Resonance: the use of two or more Lewis
structures to represent a particular molecule
or ion.
Resonance structures are connected by a
double-headed arrow. The true structure is a
blend of the resonance structures.
Resonance indicates that electron delocalization
occurs in the molecule.

25. What are the resonance structures of the carbonate (CO32-) ion?
A resonance structure is one of two or more Lewis structures for a single molecule that cannot be represented accurately by only one Lewis structure. O + - O + -
What are the resonance structures of the
carbonate (CO32-) ion? O C - O C - O C -
9.8

26. Exceptions to the Octet Rule
The Incomplete Octet
Be – 2e-
2H – 2x1e-
4e-
BeH2 H Be
B – 3e-
3F – 3x7e-
24e-
3 single bonds (3x2) = 6
9 lone pairs (9x2) = 18
Total = 24 F B
BF3
9.9

27. Exceptions to the Octet Rule
Odd-Electron Molecules
N – 5e-
O – 6e-
11e- NO N O
The Expanded Octet (central atom with principal quantum number n > 2) S F
S – 6e-
6F – 42e-
48e-
6 single bonds (6x2) = 12
18 lone pairs (18x2) = 36
Total = 48
SF6
9.9

28. How the Model Explains the Properties
of Covalent Compounds
The covalent bonding model proposes that electron sharing
leads to strong, localized bonds. But most compounds with
covalent bonds are gases, liquids, or low-melting solids.
Ionic compound are high-melting solids.
Why is this? Are covalent bonds weaker than ionic bonds?
Within a molecule there are strong intramolecular bonds.
Between molecules there are weak intermolecular forces.
The latter are what determine melting and boiling points.

29. Comparison of Ionic and Covalent Compounds
9.4

30. Properties of the Period 3 Chlorides
Fig. 9.21

31. For the test on chapter 9 study:
Test # questions 1-17
CHM Test #1 Spring 2001
Questions 1-8, 28, 29

32. Single bond < Double bond < Triple bond
The enthalpy change required to break a particular bond in one mole of gaseous molecules is the bond energy.
Bond Energy
H2 (g)
H (g) +
DH0 = kJ
Cl2 (g)
Cl (g) +
DH0 = kJ
HCl (g)
H (g) +
Cl (g)
DH0 = kJ
O2 (g)
O (g) +
DH0 = kJ O
N2 (g)
N (g) +
DH0 = kJ N
Bond Energies
Single bond < Double bond < Triple bond
9.10

34. Average bond energy in polyatomic molecules
H2O (g)
H (g) +
OH (g)
DH0 = 502 kJ
OH (g)
H (g) +
O (g)
DH0 = 427 kJ
Average OH bond energy = 2
= 464 kJ
9.10

35. Use bond energies to calculate the enthalpy change for:
H2 (g) + F2 (g) HF (g)
DH0 = SBE(reactants) – SBE(products)
Type of bonds broken
Number of bonds broken
Bond energy (kJ/mol)
Energy change (kJ) H 1
436.4 F 1
156.9
Type of bonds formed
Number of bonds formed
Bond energy (kJ/mol)
Energy change (kJ) H F 2
568.2
1136.4
DH0 = – 2 x = kJ
9.10