Basic Concepts
of Chemical
Bonding
Chapter 8
Three Types of Chemical Bonds
• Ionic bond
– Transfer of electrons
– Between metal and nonmetal ions
• Metallic bond
– Bonding electrons relatively free to move
• Covalent bond
– Sharing of electrons
– Between nonmetal atoms
Polar covalent bond ≡ a covalent bond with greater electron
density around one of the two atoms
electron poor
region
H
electron rich
region
F
e- poor
H
d+
e- rich
F
d-
Electronegativity: the ability of an atom in a molecule to attract
electrons to itself
Fig 8.6 Electronegativities of the Elements
Figure 8.7
Polar Covalent Bonds
The greater the difference in electronegativity, the more
polar is the bond.
Table 8.3
Fig 8.9
Most polar
Least polar
Writing Lewis Structures (p 314)
1. Sum up all valence electrons. Add 1 for each negative
charge. Subtract 1 for each positive charge.
2. Draw skeletal structure of compound showing what
atoms are bonded to each other. Put least
electronegative element in the center.
3. Complete octets of atoms connected to central atom.
4. Place remaining electrons on central atom.
5. If not enough electrons to give central atom an octet,
try multiple bonds.
Write the Lewis structure of nitrogen trifluoride (NF3).
Step 1 – Count valence electrons N - 5 (2s22p3) and F - 7 (2s22p5)
5 + (3 x 7) = 26 valence electrons
Step 2 – N is less electronegative than F, put N in center
Step 3 – Draw single bonds between N and F atoms and complete
octets on N and F atoms.
Step 4 - Check, are # of e- in structure equal to number of valence e- ?
3 single bonds (3x2) + 10 lone pairs (10x2) = 26 valence electrons
F
N
F
F
Write the Lewis structure of the carbonate ion (CO32-).
Step 1 – Count valence electrons C - 4 (2s22p2) and O - 6 (2s22p4)
-2 charge – 2e4 + (3 x 6) + 2 = 24 valence electrons
Step 2 – C is less electronegative than O, put C in center
Step 3 – Draw single bonds between C and O atoms and complete
octet on C and O atoms.
Step 4 - Check, are # of e- in structure equal to number of valence e- ?
3 single bonds (3x2) + 10 lone pairs (10x2) = 26 valence electrons
Step 5 - Too many electrons, form double bond and re-check # of e-
O
C
O
O
2 single bonds (2x2) = 4
1 double bond = 4
8 lone pairs (8x2) = 16
Total = 24
Formal Charges
Formal charge - 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
total number
total number
on an atom in
of valence
= electrons in - of nonbonding
a Lewis
electrons
structure
the free atom
-
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.
)
Formal Charges
• The best Lewis structure…
 …is the one with the fewest charges
 …puts a negative charge on the most electronegative atom.
Draw Lewis structure for ozone, O3
or

But this is at odds with the true, observed structure of
ozone, in which…

…both O−O bonds are the same length:
Resonance
Just as green is a synthesis of
blue and yellow…
…ozone is a synthesis of these
two resonance structures.
Resonance structure - one of two or more Lewis structures
for a single molecule that cannot be represented accurately by
only one Lewis structure.
e.g., ozone
O
O
O
O
O
O
What are the resonance structures of the
carbonate (CO32-) ion?
2-
O
C
O
O
2-
O
C
O
O
2-
O
C
O
O
Resonance
The organic compound
benzene, C6H6, has two
resonance structures:
It is commonly depicted as a hexagon with a circle inside to
signify the delocalized electrons in the ring.
Exceptions to the Octet Rule
Too few electrons
BeH2
BF3
B – 3e3F – (3)7e24e-
Be – 2e2H – (2)1e4e-
F
B
F
H
F
Be
H
3 single bonds (3x2) = 6
9 lone pairs (9x2) = 18
Total = 24
Exceptions to the Octet Rule
Odd number of electrons
NO
N – 5eO – 6e11e-
N
O
Too many electrons (central atom with principal quantum number n > 2)
SF6
S – 6e6F – 42e48e-
F
F
F
S
F
F
F
6 single bonds (6x2) = 12
18 lone pairs (18x2) = 36
Total = 48
Average Bond Enthalpies
Bond
Type
Bond
Enthalpy
(kJ/mol)
C‒C
348
C=C
614
C≡C
839
C‒N
293
C=N
615
C≡N
891
Bond Enthalpies
Single bond < Double Bond < Triple Bond
Table 8.4 Average bond Enthalpies (kJ/mol)


These are average bond enthalpies, not absolute bond enthalpies
The C−H bonds in methane, CH4, will be a bit different than
the C−H bond in chloroform, CHCl3
Chemical
Bonding
Fig 8.14 Estimating Enthalpies of Reaction
Hrxn = (bond enthalpies of bonds broken) (bond enthalpies of bonds formed)
Fig 8.14 Estimating Enthalpies of Reaction
CH4 (g) + Cl2 (g)
 CH3Cl (g) + HCl (g)
In this example:
• one C-H bond and one ClCl bond are broken
• one C-Cl and one H-Cl
bond are formed
Fig 8.14 Estimating Enthalpies of Reaction
CH4 (g) + Cl2 (g)
 CH3Cl (g) + HCl (g)
So,
H = [D(C−H) + D(Cl−Cl)] − [D(C−Cl) + D(H−Cl)]
= [(413 kJ) + (242 kJ)] - [(328 kJ) + (431 kJ)]
= (655 kJ) - (759 kJ)
= -104 kJ
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