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
Figure 9.1
A general comparison of metals and
non-metals
9-3
Types of Chemical Bonding
1. Metal with non-metal
electron transfer and ionic bonding
2. Non-metal with non-metal
electron sharing and covalent bonding
(localized)
3. Metal with metal
electron pooling and metallic bonding
(delocalized)
9-4
The three models of chemical bonding
Figure 9.2
9-5
Lewis Electron-Dot Symbols
For Main Group elements:
The 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. .
. N.
N:
:N
.
.
:
9-6
Lewis electron-dot symbols for elements in Periods 2 and 3
Figure 9.3
9-7
General Rules
For a metal, the total number of dots equals the maximum number
of electrons it loses to form a cation.
For a non-metal, the number of unpaired dots equals the number
of electrons that become paired either through electron gain or
electron sharing. The number of unpaired dots equals either the
negative charge of the anion an atom forms or the number of
covalent bonds it forms.
9-8
The Ionic Bonding Model
Involves the transfer of electrons from metal to non-metal to form
ions that come together in a solid ionic compound
The Octet Rule
When atoms bond, they lose, gain or share electrons to attain a filled
outer shell of eight (or two) electrons
In ionic bonding, the total number of electrons lost by the metal atoms
equals the total number of electrons gained by the non-metal atoms.
9-9
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 two sodiums are
needed for each oxygen.
SOLUTION:
O2-
Na
2s
O
2 Na+
.
Na
3s
3p
+ : O:
.
Na
9-10
:
2p
2Na+ + :O: 2-
:
2s
Na
2p
.
3p
.
3s
Figure 9.4
Three ways to represent the formation of Li+ and Fthrough electron transfer
1. Electron configurations
Li 1s22s1
F 1s22s22p5
+
Li+ 1s2
+
F- 1s22s22p6
2s
2p
2. Orbital diagrams
Li+
Li
1s
2s
1s
2p
+ F
+
1s
2s
F-
2p
1s
2s
+
: F: -
3. Lewis electron-dot symbols
:
9-11
Li+
:
+
:
Li .
.
:F:
2p
Ionic Bonding and Lattice Energy
The electron transfer process is an endothermic process, but ionic compound
formation is an exothermic process.
Li(g)
F(g) + eLi(g) + F(g)
Li+ + e-
IE1 = 520 kJ
F-(g)
EA = -328 kJ
Li+(g) + F-(g)
IE1 + EA = 192 kJ
But ∆Hfo for solid LiF = -617 kJ/mol!
Li+(g) + F-(g)
LiF(g)
∆Ho = -755 kJ
(an exothermic process due to the attraction of
oppositely charged ions)
9-12
Even more energy is released when the gaseous ions coalesce
into a crystalline solid. Thus….
Li+(g) + F-(g)
LiF(s)
∆Holattice of LiF = lattice energy = -1050 kJ
The lattice energy is the enthalpy change that occurs when gaseous
ions coalesce into an ionic solid.
How do we measure lattice energy experimentally? Use
Hess’s law in a Born-Haber cycle
9-13
Figure 9.6
9-14
The Born-Haber cycle for lithium fluoride
Working the Numbers
STEP 1: Enthalpy of Li atomization = 161 kJ
STEP 2: 1/2 the bond energy of F2(g)= 0.5(159 kJ) = 79.5 kJ
STEP 3: IE1 for Li(g) = 520 kJ
STEP 4: EA of F(g) = -328 kJ
The enthalpy change for the overall process, ∆Hfo, = -617 kJ
Only the lattice energy is unknown, and it is equal to the enthalpy
change of the overall process minus the sum of the above four
steps = -1050 kJ
9-15
Central Point
Ionic solids exist only because the lattice energy drives the
energetically unfavorable electron transfer.
9-16
Periodic Trends in Lattice Energy
Coulomb’s Law
charge A x charge B
electrostatic force a
distance2
But energy = force x distance. Therefore,
charge A x charge B
electrostatic energy a
distance
cation charge x anion charge
electrostatic energy a
9-17
cation radius + anion radius
a DHolattice
Trends in lattice energy
Figure 9.7
9-18
Effect of Ionic Charge on Lattice Energy
Compare LiF and MgO: Li+ and Mg2+ have similar radii, and
F- and O2- have similar radii.
∆Holattice (LiF) = -1050 kJ/mol
∆Holattice (MgO) = -3923 kJ/mol
The nearly four-fold larger value for MgO reflects the difference in
the product of the charges (12 vs 22) in the numerator of the
electrostatic energy equation (monovalent vs divalent ions).
9-19
Does the ionic model explain the
properties of ionic compounds?
9-20
Electrostatic forces
and the reason ionic
compounds crack
9-21
Figure 9.8
Electrical
Conductance
and Ion Mobility
Figure 9.9
Solid ionic
compound
9-22
Molten ionic
compound
Ionic compound
dissolved in water
Table 9.1
Melting and Boiling Points of Some Ionic Compounds
Compound
bp (oC)
CsBr
636
1300
NaI
661
1304
MgCl2
714
1412
KBr
734
1435
CaCl2
782
>1600
NaCl
801
1413
LiF
845
1676
KF
858
1505
2852
3600
MgO
9-23
mp (oC)
Vaporizing an
ionic compound
Figure 9.10
9-24
The Covalent Bonding Model
Each atom in a covalent bond “counts” the shared electrons
as belonging entirely to itself.
An electron pair that is part of an atom’s valence shell but not
involved in bonding is called a lone pair, or unshared pair.
Bond order: the number of electron pairs being shared between
any two bonded atoms
single bond (H2) - bond order of 1
double bond (H2C=CH2) - bond order of 2
triple bond (N2) - bond order of 3
9-25
Covalent bond
formation in H2
Figure 9.11
9-26
The attractive and repulsive
forces in covalent bonding
Figure 9.12
9-27
Properties of Covalent Bonds
Bond energy (bond enthalpy or bond strength): the energy required
to overcome the mutual attraction between the bonded nuclei and
the shared electrons.
Bond breakage is an endothermic process; bond energy is always
positive.
Bond formation is an exothermic process.
9-28
Bond Length
For a given pair of atoms, a higher bond order results in a
shorter bond length and a higher bond energy.
A shorter bond is a stronger bond.
9-29
9-30
9-31
Bond length and covalent radius
internuclear distance
(bond length)
internuclear distance
(bond length)
9-32
covalent
radius
covalent
radius
internuclear distance
(bond length)
internuclear distance
(bond length)
Figure 9.13
covalent
radius
covalent
radius
9-33
SAMPLE PROBLEM 9.2
PROBLEM:
Comparing Bond Length and Bond Strength
Using the periodic table, 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) Bond order =1 for all and sulfur is bonded to halogens; bond
length should increase and bond strength should decrease with
increasing atomic radius. (b) Similar atoms (C) are bonded but
bond order changes; bond length decreases as bond order
increases, and bond strength increases as bond order increases.
SOLUTION:
(a) Atomic size increases
moving down a group.
(b) Using bond orders we get:
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
9-34
O
O>C=O>C-O
Properties of Covalent Compounds
Weak forces between molecules, not the strong covalent
bonds within each molecule, are responsible for the
physical properties of covalent compounds.
Covalent compounds have relatively low melting and boiling
points.
Most covalent compounds are poor electrical conductors.
9-35
Strong forces within molecules, weak forces between them
Figure 9.14
Strong covalent bonding forces within molecules
Weak intermolecular forces between molecules
9-36
Network Covalent Solids
No separate molecules; held together by covalent bonds that
extend throughout the sample
quartz: melts at 1550 oC.
diamond: melts at 3550 oC.
These examples illustrate the strength of covalent bonds.
9-37
Covalent bonds of network covalent solids
Figure 9.15
9-38
The Concept of Electronegativity (EN)
Defined as the relative ability of a bonded atom to attract
shared electrons (not the same as EA)
Bond energy of H2 = 432 kJ/mol
Bond energy of F2 = 159 kJ/mol
Bond energy of HF = 565 kJ/mol, not 296 kJ/mol
The stronger-than-expected HF bond is due to unequal
sharing of electrons, with F bearing a partial negative
charge and H bearing a partial positive charge. The
attraction between the partial charges strengthens the
bond.
9-39
The Pauling electronegativity (EN) scale
Figure 9.16
9-40
Trends in Electronegativity
In general, electronegativity is inversely related to
atomic size.
For main-group elements, EN generally increases up
a group and across a period.
Non-metals are more electronegative than metals.
The least electronegative (most electropositive) nonradioactive element is Cs (lower left-hand corner of
the Periodic Table).
9-41
Electronegativity and atomic size
Figure 9.17
9-42
Electronegativity and Oxidation Number
(a) The more electronegative atom in a bond is assigned
all of the shared electrons; the less electronegative atom is
assigned none of the shared electrons.
(b) Each atom in a bond is assigned all of its unshared
electrons.
(c) The oxidation number is given by:
O.N. = # valence e- - (# shared e- + # unshared e-)
e.g.: HCl: Cl more electronegative than H; has 7 valence
electrons; has an O.N. of 7 - 8 = -1
H has 1 valence electron; has an O.N. of 1 - 0 = +1
9-43
Polar Covalent Bonds and Bond Polarity
Covalent bonds involving atoms with different electronegativities:
generate partial (+) and (-) charges; defined as polar covalent
bonds (e.g., HCl)
Polar covalent bonds: depicted by a polar arrow (
points toward the negative pole
H2 and F2: examples of nonpolar covalent bonds
9-44
) that
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 to find EN values; the arrow should point
toward the negative end.
(b) EN 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
9-45
Partial Ionic Character of Polar Covalent Bonds
Related directly to the electronegativity difference (∆EN) between
the bonded atoms
The greater the ∆EN, the larger the partial charges and the higher
the partial ionic character (PIC).
Thus LiF has more PIC than HF; HF has more PIC than F2.
9-46
3.0
DEN
2.0
Boundary ranges for
classifying the ionic
character of chemical
bonds
Figure 9.18
9-47
0.0
Percent ionic character as a function of
electronegativity difference (DEN)
9-48
Figure 9.19
Charge density of
the LiF molecule
(an ionic compound)
Li
F
No bond has 100%
ionic character; electron
sharing occurs to some extent
Figure 9.20
9-49
Ionic-To-Covalent Bonding Continuum Across a Period
Consider bonding between a metal and non-metal in Period 3
NaCl, MgCl2, AlCl3, SiCl4, PCl3, S2Cl2, and Cl2
Increasing covalent character (decreasing ionic character)
from NaCl to Cl2
Underlying factor: As ∆EN becomes smaller, the bond becomes
more covalent.
9-50
Figure 9.21
9-51
Properties of the Period 3 chlorides
Metallic Bonding
The electron-sea model: all metal atoms in the sample
contribute their valence electrons to form an “electron
sea” that is delocalized throughout the substance
The metal atoms are not held in place as rigidly as are
the ions of an ionic solid.
9-52
Table 9.5
9-53
Melting and Boiling Points of Some Metals
element
mp (oC)
bp (oC)
lithium (Li)
180
1347
tin (Sn)
232
2623
aluminum (Al)
660
2467
barium (Ba)
727
1850
silver (Ag)
961
2155
copper (Cu)
1083
2570
uranium (U)
1130
3930
Melting points of the Group 1A and Group 2A elements
Figure 9.23
9-54
The reason metals deform
Figure 9.24
9-55
metal is deformed
Tools of the Laboratory
Infrared Spectroscopy
Figure B9.1
Some vibrational modes in general diatomic and triatomic molecules
9-56
Tools of the Laboratory
Infrared
Spectroscopy
Some vibrational
modes in general
diatomic and triatomic
molecules
Figure B9.1
9-57
Tools of the Laboratory
Some vibrational modes
in general diatomic and
triatomic molecules.
Figure B9.1
9-58
Infrared Spectroscopy
Tools of the Laboratory
Figure B9.2
The infrared (IR) spectrum of acrylonitrile
9-59