Chapter 8
Chemical Bonds
The block being supported by
soap bubbles is a piece of
SEA gel, a solid foam made
from agar. Aerogels like SEA
gel have a porous structure
that encapsulates gases. It is
claimed that the density of an
aerogel can be equal to or
less than that of air. This
material, which was
developed for the space
program, is a good insulator
and can be used in
packaging. The properties of
SEA gel and other materials
are related to the types of
bonds between their atoms.
In this chapter we examine
the formation of bonds.
Assignment for Chapter 8
8.46, 8.50,8.51, 8.56, 8.60, 8.65, 8.68, 8.71
How atoms are linked together?
A+BAB
EA+EB>EAB
Generally,
An+Bm+Ck+…AnBmCk…
Ionic bonds
Valence bonds
Figure 8.1
Gilbert Newton Lewis (1875–1946).
Lewis Symbols for Atoms and Ions
H
He
N
O
K Mg
duplet
K + Cl K [ Cl ]
octet
Nitrogen atomic orbitals
N
Lewis symbol
Figure 8.2 When a main-group metal atom forms a cation, it loses its
valence s- and p-electrons and acquires the electron configuration of
the preceding noble-gas atom. The heavier atoms in Groups 13 and
14 behave similarly, but the resulting core consists of the noble-gas
configuration and an additional complete subshell of d-electrons.
1s2
+2s2 2p6
+3s2 3p6
+4s2 4p6 3d10
Li+ Be2+
Na+
B3+
Mg2+
K+
Ga3+
Rb+ In3+/5+
+5s2 5p6 4d10
Cs+ Tl3+/5+
+6s2 6p6 4f14 5d10
Fr+ Ra2+
Figure 8.3 When p-block elements acquire electrons and form anions,
they do so until they have reached the electron configuration of the
following noble gas.
1s2
H-
+2s2 2p6
O2-
+3s2 3p6
P3Se2-
Te2-
+4s2 4p6 3d10
+5s2 5p6 4d10
Bi3-
+6s2 6p6 4f14 5d10
Ionic Bonds
• An ionic bond is the electric attraction
between the opposite charges of
cations and anions.
+
-
The formation of ionic bonds is represented in terms of Lewis
symbols by the loss or gain of electrons until both species have
Reached an octet or duplet of electrons.
Figure 8.4 When tin(II) oxide is heated in air, it becomes
incandescent as it reacts to form tin(IV) oxide. Even without being
heated, it smolders and can ignite.
SnO + O2 SnO2
2- 4+
2+
22Sn [ O ] + O [ O ]Sn[ O ]
Lewis symbols and formulae are not
really written for transition metals.
Variable Valence
More than one type of cation can be formed for some elements:
Inert-pair effect
The tendency to form cations two units lower in
charge than expected from the group number, i.e.,
the p-block elements to form cations that have octet
configuration.
• 4p,5p,6p
• 4s,5s,6s
Figure 8.5 The typical ions formed by the heavy elements in Groups
13–15 show the influence of the inert-pair effect. These elements
have the tendency to form compounds in which the oxidation
numbers differ by 2.
3+
。。
。。
。。
InCl3: In [:C l :][:C l :][:C l :]
。。
。。
。。
+
。。
InCl: In [: C l :]
。。
Figure 8.6 Considerable energy is needed to produce cations and
anions from neutral gas-phase atoms: the ionization energy of the
metal atoms must be supplied, and it is only partly recovered from
the electron affinity of the nonmetal atoms. The overall lowering of
energy that drags the ionic solid into existence is due to the strong
attraction between cations and anions in the solid. It takes 145 kJ/mol
to produce the ions from the elements, and the solid compound is
787 kJ/mol lower in energy than the separated ions.
Figure 8.7 A two-dimensional slice of an ionic crystal. The greater the
distance between two ions, the weaker their attractions. Therefore, the
strongest attractions to an ion are those of the adjacent ions of opposite
charge. The blue circles denote the distances over which the two closest
repulsions occur, and the yellow circles denote the two closest
attractions.
Vij 
qi q j
rij
N N
N N
N N
j 1 i 1
j 1 i 1
j 1 i 1
E  12 [ | Vij |   | Vij |   | Vij |]
atomization
element(substance) 
 atoms
ionization

 ions
formation
element(substance) 
 solid
lattice formation
ions 
 solid
lattice enthalpy =enthalpy of atomization
+enthalpy of ionization - enthalpy of formation
Figure 8.8 In a Born-Haber cycle, we select a
sequence of steps that starts and ends at the
same point (the elements, for instance). The
lattice enthalpy is the enthalpy change of the
step in which the solid is formed from a gas of
ions (the dark red arrow). The sum of enthalpy
changes for the complete cycle (yellow line) is
0 because enthalpy is a state property. The
“Enthalpy of ionization” step includes the
energies required to produce both cations and
anions from their respective atoms.
Figure 8.9 The Born-Haber cycle used to determine the lattice
enthalpy of potassium chloride.
Enthalpy of ionization (Cl)
Enthalpy of ionization (K)
Enthalpy of atomizaiton (Cl)
Enthalpy of atomizaiton (K)
-Enthalpy of formation
Lattice enthalpy:
H L (KCl)  89  122  418  437  349
 717 kJ/mol
Figure 8.10 An ionic solid—here we see a fragment of sodium
chloride, with the sodium ions represented by pink spheres and the
chloride ions by green spheres—consists of an almost infinite array
of cations and anions stacked together to give the lowest energy
arrangement. The pattern shown here is repeated throughout the
crystal.
Figure 8.11 This sequence of images illustrates why ionic solids are brittle. (a)
The original solid consists of an orderly array of cations and anions. (b) A
hammer blow can push the ions into positions where cations are next to
cations and anions are next to anions; there are now strong repulsive forces
acting (as depicted by the double-headed arrows). (c) As a result of these
repulsive forces, the solid springs apart in fragments. (d) This chunk of calcite
consists of several large crystals joined together. (e) The blow of a hammer
has shattered the crystal, leaving flat, regular surfaces consisting of planes of
ions.
Figure 8.12 These micrographs show the porous structure of bone.
The calcium in bone is extracted by the body if the level of calcium in
the diet is low. (a) Healthy bone tissue. (b) Bone that has suffered
calcium loss through osteoporosis. The overlay shows the regular
arrangement of the calcium and phosphate ions in healthy bone.
(a)
(b)
Covalent Bonds
• In covalent bond formation, atoms go as far as
possible toward completing their octets (duplets) by
sharing electron pairs.
Figure 8.13 The formation of a covalent bond between two hydrogen
atoms. (a) Two separate hydrogen atoms, each with one electron. (b)
The electron cloud that forms when the spins pair and the orbitals
merge is most dense between the two nuclei. (c) The boundary
surface that we shall use to depict a covalent bond.
Covalent Bonds
• In covalent bond formation, atoms go as far as
possible toward completing their octets
(duplets) by sharing electron pairs.
H+H  H:H or H-H
..
:
..
..
..

:
. :
:
: or
F F FF
.
..
..
..
That’s why F2 is so reactive.
:
..
..
..
..
..

:
F F
Lone pairs
The Lewis structure gives not only bond positions (shared electron pairs)
But also reactivity-how reactive is the molecules and which part(s) is(are)
reactive (lone pairs).
The Structures of Ployatomic Species
H

H:C:H

H
H
|
H- C -H
|
H






:F O F:

:O:
H
H
|
|
|


[H- N -H]+[:O S  O:]2-[H- N -H]+
|


|
|
:O:
H
H

• Single, double, triple bonds may be
formed.
How to write Lewis structure of a
polyatomic species
• Count the total number of valence electrons on each atom
and divide by 2 to obtain the number of electron pairs. If the
species is polyatomic ion, add one more electron for each
negative charge or subtract one electron for each positive
charge.
• Write the chemical symbols of the atoms to show their
layout in the molecule. One can predict the most likely
arrangements of atoms by using common patterns and the
rules of thumb (largely correct but exceptions possible)
given earlier.
• Place one electron pair between each pair of bonded
atoms.
• Complete the octet (duplet for H) of each atom by placing
any remaining electron pairs as lone pairs around the
atoms. If there are not enough electrons to form octets,
form multiple bonds.
More Examples
NH3:
|
3+5=8 valence electrons = 4 pairs.
H- N :
|
3 pairs to form three bonds,
H
leaving one lone pair
H




H  O Br :
Hypobromous acid HBrO:
1+6+7=14 valence electrons = 7 pairs
Two pairs to form two bonds,
Leaving 5 lone pairs.
:O:
||
H- C
|
: O -H

Formic acid HCOOH
2+12+4=18 valence electrons
= 9 pairs
5 pairs to form five bonds,
leaving 4 lone pairs.
Classroom Exercise
• Write the Lewis structure of CH3COOH
H
:O:
|
||
H- C C
|
|
H : O -H

Resonance: Blending of Structures

:O
|| 

[:O N  O:]


:O:
| 

[:O  N  O:]

:O:
| 

[:O N  O:]-
All valid. We cannot find two bond lengths (hypothetical N-O vs N=O)
Resonance: Blending of Structures
Resonance: Blending of Structures
Quiz
Write the Lewis structure of the following compounds:
(1) Methanol (2) Nitrate (3) Carbon Monoxide
Answer:
:O:
||
H- C
|
: O-H


:O
|| 

[:O N  O:]


:O:
| 

[:O  N  O:]

:O:
| 

[:O N  O:]
Formal Charge
• The real number of valence electrons each atom in a molecule
“owns” may be different from the number of normal valence
electrons of that atom, rendering it positively or negatively charged.
Formal charge (FC) = number of valence electrons on the free atom (V)
- number of electrons present as lone pairs (L)
- ½ number of electrons shared in bonds (S/2)
FC  V  L  S
1
2
Formal charge can be understood as surplus or deficit of valence electrons after
deducting lone pairs and shared pairs.
Typically, the most stable Lewis structures are those in which the formal charge
of the individual nonmetal atoms are close to zero
Plausibility of a structure






OCO
NO N
0
-1 +2 -1


0

0

OOC

0 +2 -2

0 +2 -2




N NO
-1 +1 0
Classroom Exercise
• Suggest a plausible structure for the poisonous
gas phosgene, COCl2.
Resonance
states

:Cl:
0
|

O C

|
:Cl:

0
0
0
0
..O
:Cl:
 ||
 |
:Cl C :Cl C
.. |
.. ||
:Cl:
:O


0
  ..
:Cl-O - C 0
..  |
:Cl: 0

Formal charges are all zero, too.
But why is this one not favored?
Experiment shows the oxygen has
double bond rather than single bond.
Exceptions to the Octet Rule
•
•
•
•
B, C, N, O, F observe the octet rule.
P, S, Cl can have more than 8 valence electrons.
Radicals: odd-electron species, highly reactive.
Biradicals: with two unpaired electrons.
CH3 -CH3 
 H3C  +  CH 3
stress
O2 (g)  2O(g)
sunlight
O+O 2  O3  O+O 2
UV








O  O O: :O O  O


O2 is a biradical!


OO
Wrong!




O O

..

..
.
..
O
..
.. 2
.
Classroom Exercise
Write a Lewis structure for the hydrogenperoxyl radical HOO.,
which plays an important role in atmospheric chemistry and
which in the body has been implicated in the degeneration of
neurons.
Case Study 8 (a) The equipment in the illustration monitors air quality
from a rooftop in Los Angeles. The concentration of NO2 in the air
due to automobile traffic increases during the day, contributing to the
typical brown color of the afternoon sky.
Exhaust: N 2 (g)+O 2 (g)  2NO(g)
+Atmosphere: 2NO(g)+O 2 (g)  2NO 2 (g)
UV (< 400 nm)
+Sunlight:  NO 2 
 NO 2 +  O  (smog)
Case Study 8 (b) The gas NO (left) is colorless. When it is exposed to
air (right), it is rapidly oxidized to brown NO2.
Oxygen molecules are
actually biradicals.
2NO(g)+O2 (g)  2NO2 (g)
Expanded valence shells
• Nonmetal atoms in Period 3 or higher can
accommodate 10, 12 or more electrons (expanded
valence shell). e.g., P, S, Cl.
P 4 (g)+6Cl2 (g)  PCl3 (l)
limited supply of chlorine
excessive chlorine
P 4 (g)+10Cl2 (g) 
 4PCl5 (s)
Figure 8.16 (a) A model using small spheres to
represent atoms and (b) a space-filling model of
PCl5, showing how closely the chlorine atoms must
pack around the central phosphorus atom. (c) A
nitrogen atom is significantly smaller than a
phosphorus atom, and five chlorine atoms cannot
pack around it.
Figure 8.17 Phosphorus trichloride is a colorless
liquid. When it reacts with chlorine (the pale yellowgreen gas in the flask), it forms the very pale yellow
solid phosphorus pentachloride (at the bottom of
the flask).
+
Figure 8.18 Two circumstances in which a central atom
assumes an expanded octet. First, there are too many
electrons to be accommodated in octets because either (a)
there are too many atoms attached, or (b) the central atom
must accommodate additional lone pairs. (c) Second, a
resonance structure with multiple bonds has a favorable
energy.
S accommodates 10 electrons
How many electrons does Xe
accommodate?
Which is most plausible?
12 valence electrons!
Classroom exercise:
which is more plausible?
Answer
One more…
The unusual structures of group
13 halides
• Compounds of boron and aluminum may
have unusual Lewis structures in which
boron and aluminum have incomplete
octets or in which halogen atoms as
bridges.
Incomplete octet
+
F
Coordinate covalent bond
This covalent bond is
formed by two electrons
from nitrogen.
BF3 (g)+NH3 (g)  NH3BF3 (s)
Both electrons shared in the bond come from the same atom.
The coordinate bond is more extended (nonlocal) than an ordinary covalent bond.
Coordinate covalent bond
Which atom contributes the electrons that form the coordinate covalent bond?
AlCl3 ( s) 
 Al2Cl6 (g)
sublimes at 180 o C
>200 o C
Which atom contributes the electrons that form the coordinate covalent bond?
AlCl3 ( s) 
 Al2Cl6 (g)
sublimes at 180 o C
>200 o C
Lewis Acids and Bases
• Lewis acid: electron pair acceptor
• Lewis base: electron pair donor
• Lewis acid-base complex: product of
reaction between a Lewis acid and a
Lewis base:
acid + :base  complex
+
Acid
F
Base
Complex
HCl  H + +ClH
H +
|
|
+:O -H  H  O -H
..
..
CaO+H 2 O  Ca(OH) 2
H
|
..
..
[:O:]2- + O -H  2[:O -H]..
..
H
..
+
base
complex
acid
base
H
H
|
|
|
..
+
H- N :  O -H  [H- N -H] +[:O -H]|
|
..
..
H+
H
H
H
base
complex
acid
acid
base
Figure 8.19 Oxygen molecules are transported through our
bloodstream in the form of a complex with the iron atoms in
hemoglobin molecules and then released where they are needed.
Carbon monoxide forms stronger
with iron, causing oxygen starvation.
biradical
complex
acid
Cations of d-block metals are often good Lewis acids.
Quiz
1. Find which of the following structures is more plausible based on formal charges:




N NO




NO N
2. Write the correct Lewis structure of oxygen molecules.
3. Identify the Lewis acid and Lewis base in the following reaction:
CaO+H 2 O  Ca(OH) 2
H
|
..
..
[:O:]2- + O -H  2[:O -H]..
..
..
Ionic vs Covalent Bonds
Ionic bonds: electron gain or loss.
Covalent bonds: electron pair(s) shared
equally.
Ionic and covalent bonds are two
extreme cases in explaining the
formation of molecules from atoms.
Most actual chemical bonds are in the
between.
Figure 8.20 The electronegativity of an element is “its electron-pulling power
when it is part of a compound.” (a) An atom (B) with a high electronegativity has
a strong pulling power on electrons (as represented by the large arrow),
particularly for the electron pair it shares with its neighbor. (b) The outcome of
the tug-of-war: the more electronegative atom has a greater share in the
electron pair of the covalent bond.
Electronegativity is a description of “unequally
sharing of electron pair(s)”.
Figure 8.21 The variation of the electronegativity of the main-group
elements (with the exception of the noble gases). Electronegativity
tends to be high toward the upper right corner of the periodic table
and low on the lower left. Elements with low electronegativities,
such as the s -block metals, are often called electropositive.
Figure 8.22 The electronegativity difference between two elements can be
used to predict the most appropriate bonding model for a chemical bond
between the elements. When the difference is small (less than about 1.5), the
covalent model is good. When the difference is large (greater than about 2), it
is most accurate to express the bonding in terms of the ionic model. Note
that the numbers we have quoted are only a guide.
Figure 8.23 When a small, highly charged cation is close to a
large anion, the electron cloud of the latter is distorted in the
process we call polarization. Small, highly charged cations are
highly polarizing. Large, electron-rich anions are highly
polarizable.
Compounds composed of highly polarizing cations and highly
polarizable anions have a significant covalent character in their
bondings.
Polarizability is a description of “ transition from electron gain to electron
sharing.”
Figure 8.24 The polarizability of an anion and the polarizing power of
a cation can be used as a guide to judge whether an ionic model of a
bond is likely to be valid. As the polarizing power and polarizability
increase, the distortion of the anion becomes so great that a covalent
description of the bond becomes more reasonable.
Figure 8.25 The silver halides, which were formed here by
precipitation from silver nitrate and sodium halide solutions,
become increasingly insoluble down the group from AgCl to AgI, as
the polarizability of the anion increases. The figure shows, from left
to right, AgCl, AgBr, and AgI.
Covalent character increases.
Be-Cl bond is significantly covalent
and long chains form in the solid.
Cl
Be
Cl-2? And Be+4?
Yes and No.
詩云
Lewis 模型頌
Ode to the Lewis Model
價層電子傾二八
成鍵孤對皆可納
形式電荷辯真偽
鹵素磷硫配橋差
Valence electrons tend to form octet or duet,
Where either bonded or lone pairs are accepted.
Formal charges tell the correct from incorrect,
But halogens, P, S, coordinated and bridged atoms may be except.
Assignment for Chapter 8
8.46, 8.50,8.51, 8.56, 8.60, 8.65, 8.68, 8.71