Chapter 9
Models of Chemical
Bonding
9-1
Models of Chemical Bonding
9.1 Atomic Properties and Chemical Bonds
9.2 The Ionic Bonding Model
9.3 The Covalent Bonding Model
9.4 Between the Extremes: Electronegativity and Bond Polarity
9.5 An Introduction to Metallic Bonding
9-2
A general comparison of metals and
nonmetals
9-3
Types of Chemical Bonding
1. Metal with nonmetal:
electron transfer and ionic bonding
2. Nonmetal with nonmetal:
electron sharing and covalent bonding
3. Metal with metal:
electron pooling and metallic bonding
9-4
The three models of chemical bonding
9-5
Lewis Electron-Dot Symbols
For main group elements The A group number gives the number of valence electrons.
Place one dot per valence electron on each of the four
sides of the element symbol.
Pair the dots (electrons) until all of the valence electrons are
used.
Example:
Nitrogen, N, is in Group 5A and therefore has 5 valence
electrons.
.
: N.
.
9-6
Lewis electron-dot symbols for elements in Periods 2 and 3
9-7
Ionic Bonding
Ionic bonding results when there is a transfer
of electrons between two atoms. Each atom
achieves a full outer level of electrons.
For many atoms in the 2nd and 3rd period, this
would be 8 electrons, known as the octet rule.
9-8
SAMPLE PROBLEM 9.1
PROBLEM:
PLAN:
Depicting Ion Formation
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.
Draw orbital diagrams for the atoms and then move electrons to
make filled outer levels. It can be seen that 2 sodiums are
needed for each oxygen.
SOLUTION:
O2-
Na
2s
O
2 Na+
.
Na
3s
3p
:
2p
+ : O:
.
Na
2Na+ +: O:2-
:
2s
Na
2p
.
3p
.
3s
Na2O
9-9
Three ways to represent the formation of Li+ and Fthrough electron transfer.
Electron configurations
Li 1s22s1
F 1s22s22p5
+
Li+ 1s2
+
F- 1s22s22p6
2s
2p
Orbital diagrams
Li+
Li
1s
2s
1s
2p
+
+ F
1s
2s
F1s
2p
2s
Lewis electron-dot symbols
:
9-10
Li+
+
: F: -
:
+
:
Li .
.
:F:
2p
Born-Haber Cycle
Lattice Energy - Energy released when ions come together
forming an ionic solid.
Remember Hess’s law states that the enthalpy change between
two states is the same as the sum of enthalpies in a multistep
process that goes between the same two states.
H0f =  H0 elements to atoms +  H0 ions from atoms + H0lattice
So from this one equation if all of the H0 ‘s except one are known,
e.g. H0lattice , the value can be calculated.
9-11
The Born-Haber cycle for lithium fluoride
9-12
Trends in lattice energy
9-13
Covalent Bonding
Covalent bonding results when two electrons are shared in
an orbital between two atoms. Each atom achieves a full
outer level of electrons resulting in a lower energy system.
The pair of electrons used are called the shared or bonding
pair. In terms of the octet rule, this pair of electrons counts
for both atoms in completing the octet.
The electron pairs that are not involved in bonding belong
only to the atom with which they are associated. These are
called lone pairs.
BOND ORDER - When only one pair of electrons are shared
between two atoms, it’s called a single bond.
If two pairs of electrons are shared covalently between two
atoms, it’s called a double bond; three pairs, triple bond.
9-14
Covalent bond formation in H2.
9-15
The attractive and repulsive
forces in covalent bonding.
9-16
Bond Energy - The amount of energy required to break a bond.
The greater the energy, the stronger the bond.
Bond breaking is an endothermic process, so bond breaking
enthalpies are positive.
9-17
Bond Length - In general, the closer the electrons are
held by the atoms, the shorter the bond length and
the higher the bond energy.
Multiple bonds result in stronger, shorter bonds.
9-18
SAMPLE PROBLEM 9.2
PROBLEM:
Comparing Bond Length and Bond Strength
Using the periodic table, but not Tables 9.2 and 9.3, rank the
bonds in each set in order of decreasing bond length and bond
strength:
(a) S - F, S - Br, S - Cl
PLAN:
(b) C = O, C - O, C
O
(a) The bond order is one for all and sulfur is bonded to halogens;
bond length should increase and bond strength should decrease
with increasing atomic radius. (b) The same two atoms are
bonded but the bond order changes; bond length decreases as
bond order increases while bond strength increases as bond order
increases.
SOLUTION:
(a) Atomic size increases
going down a group.
Bond length: S - Br > S - Cl > S - F
Bond length: C - O > C = O > C
Bond strength: S - F > S - Cl > S - Br
Bond strength: C
9-19
(b) Using bond orders we get
O
O>C=O>C-O
Electronegativity and Bond Polarity
Electronegativity, (EN), is the ability of an atom to attract electron
density of shared electrons.
To the extent that an atom attracts extra electron density away
from the other atom, it has a partial negative charge. The other
atom has a corresponding positive charge. This creates a polar
bond. The greater the difference in EN between the two atoms,
the more polar the bond.
The extreme “polar bond” is where the electrons are completely
with one atom from the other. What has already been identified as
an ionic bond.
On the Pauling scale of EN, fluorine is the most electronegative
with a value of 4.0. All other elements are less electronegative
with Cesium being the least with a 0.7 value.
9-20
The Pauling electronegativity (EN) scale.
9-21
Electronegativity and atomic size.
9-22
SAMPLE PROBLEM 9.3
PROBLEM:
Determining Bond Polarity from EN Values
(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) Use Figure 9.16(button at right) to find EN values; the
arrow should point toward the negative end.
(b) Polarity increases across a period.
SOLUTION: (a) The EN of N = 3.0, H = 2.1; F = 4.0; I = 2.5, Cl = 3.0
N-H
F-N
I - Cl
(b) The order of increasing EN is C < N < O; all have an EN
larger than that of H.
H-C < H-N < H-O
9-23
3.0
EN
2.0
Boundary ranges for
classifying ionic character
of chemical bonds.
0.0
9-24
Percent ionic character of electronegativity difference (EN).
9-25
9-26
Properties of the Period 3 chlorides.
End of Chapter 9
9-27