CHAPTER 8
MODELS OF CHEMICAL BONDING
Recall Section 2.5: Chemical bonds hold atoms together in a compound.
 transfer of electrons, forming cations and anions, results in ionic bonding
 sharing of electron pairs results in covalent bonding
One additional type of bonding: sharing of bulk quantity of electrons by many atoms: metallic bonding
All three types of bonding are electrostatic (attraction between + and – charges) at their hearts.
~Table 8.1
Lewis Dot Symbols
A way to depict valence electrons.
e.g., for S: [Ne]3s 3p for a total of 6 valence electrons [S is in Group 6A(16)]
2
4
The Octet Rule
ATOMS TEND TO GAIN, LOSE, or SHARE ELECTRONS to ATTAIN A FILLED OUTER SHELL of 8 ELECTRONS.
8-1
Covalent Bonding
A covalent bond consists of a pair of electrons shared by two atoms.
energy diagram of H2 bond formation
balance of attractions & repulsions
bond length: distance between nuclei of bonded atoms. Reported as averages just like bond enthalpies.
bond order: number of electron pairs shared between two bonded atoms
BO = 1
BO = 2
BO = 3
For a given pair of atoms:
The HIGHER the bond order, the SHORTER the bond, and the HIGHER the bond energy.
8-2
Properties of Covalent Compounds
Covalent bonds WITHIN molecules: STRONG
Intermolecular forces BETWEEN molecules: WEAK
Covalent molecular substances:
 Poor electrical conductors (even molten or aqueous)
 Relatively low melting points
Quartz, graphite, diamond
 are examples of “network covalent solids”.
 don’t exist as individual molecules.
Electronegativity
Electronegativity expresses the ability of a bonded atom to attract electrons to itself.
Pauling scale: F assigned EN = 4.0
PROTIP
EN and Oxidation Number
A useful tool for determining ON for atoms where Chapter 4 rules don’t help.
1. More electronegative atom assigned all shared electrons
2. Lone (unshared) pairs of electrons assigned to owning atom
3. ON = (valence e– in neutral atom) – (total e– assigned)
e.g., cyanide ion CN–
8-3
Being the most electronegative, F
always has ON of –1 in
compounds. Also, we can now see
why H has ON +1 with nonmetals
and –1 with metals, based on its
EN of 2.1 (lower than that of other
nonmetals; higher than that of
metals and boron). In a bond, the
more electronegative atom carries
the more negative ON.
Bond Polarity
Unequal sharing of electrons in a covalent bond: “polar covalent bond”
e.g., H–F
ΔEN tells us how ionic or covalent a bond is.
8-4
Depicting Molecules & Ions with Lewis Structures
Steps for drawing Lewis structures
1. Sum the valence electrons from all atoms. For anions, add one electron for each negative charge.
For cations, subtract one electron for each positive charge.
2. Write the symbols for the atoms to show which are connected to which. Connect the atoms with
single bonds. The least electronegative element is usually the central atom. This means that in
general, the atom with the lower group number is the central atom. If all atoms are in the same
group, the atom lower in the group is the central atom.
3. For each bond in the skeletal structure from step 1, subtract 2 electrons from the total.
4. Complete octets of atoms bonded to the central atom. Exception: H only can have 2 electrons.
5. Place any leftover electrons on the central atom, even if doing so results in the central atom
having more than an octet.
6. If there are not enough electrons to give all atoms octets, start using multiple bonds.
Examples
 Draw the Lewis structure for sulfur tetrafluoride, SF4.
1. 1(S) + 4(F) = 1(6) + 4(7) = 34 valence electrons
F
F
F
F
F
F
2.

S
4.
S
F
F
F
F
(32 e–)
F
5.
Draw the Lewis structure for carbon dioxide, CO2.
1. 1(C) + 2(O) = 1(4) + 2(6) = 16 valence electrons
4.
6.

S
O
O
C
O
C
O (16 e–)
2. O
5. no extra electrons to put on C.
Draw the Lewis structure for nitrogen triiodide, NI3.
8-5
C
O
F
(34 e–)
Resonance Structures
 Equivalent Lewis structures that differ in the placement of the same number/types of bonds.
 Draw the Lewis structure for sulfur dioxide, SO2.
S
O
1. 1(S) + 2(O) = 1(6) + 2(6) = 18 valence electrons
2. O
4.
6.
O
O
S
O
(16 electrons used)
S
O
↔
O
S
O
5.
O
The S-O bond order would be
S
O
(all 18 electrons used up)
(two resonance structures of SO2)
3 electron pairs
3
1
= = 12
2 atom-to-atom linkages 2
Finding the “Best” Structure When There Is a Real Choice

Draw the Lewis structure for thiocyanate ion, SCN–.
1. 1(S) + 1(C) + 1(N) + 1(–1 charge) = 6 + 4 + 5 + 1 = 16 valence electrons
2. S
C
N
S
C
N
S
C
S
N
C
N
???
or
or
…6.
–
Which Lewis structure of SCN to choose? Calculate formal charge on each atom in each structure.
Formal Charge
Rules for calculating formal charge:
1. All unshared electrons on an atom are assigned to the atom on which they are found.
2. Half of the bonding electrons are assigned to each atom in a bond.
3. FC = (number of valence e– in isolated atom) – (e– assigned)
The preferred (major) Lewis structure will have (1) the smallest magnitude FCs and (2) any negative FCs
on atoms with higher electronegativities. Therefore:
S
Formal Charge
+1
C
0
N
S
C
N
S
C
N
–2
0
0
–1
–1
0
0
The middle Lewis structure has the smallest FCs and places the negative FC on nitrogen, which is the
most electronegative atom of the three. Therefore, the center structure (S=C=N–) is preferred if we must
pick only one. It is the major contributor to the structure. The right-hand structure is a minor contributor.
8-6
Exceptions to the Octet Rule
1. Odd Number of Electrons: “free radicals”
e.g., ClO2
Possibly obeyed more
than the octet rule
2. Electron-Deficient Molecules (primarily B, Be central atoms)
e.g., BeF2
e.g., BF3
3. Expanded Valence Shells
Seeming use of empty valence d orbitals in addition to s and p orbitals
8-7