Chapter 3.1
Structures and Properties of
Substances
Chemical Bonding
The orbitals in the Periodic Table
highly
elements,
moderately
andis unreactive
The elements of the periodic table
canreactive
be classified
according
to thereactive
type of elements,
orbital that
being filled.
elements. While most main group elements are solids at room temperathe roughly
s blockone
and
the of
p them
blockareare
called
either
the main group
quarter
gases,
and one
is a liquid.
• Elements that appear in ture,
elements or the representative
elements.
Elements
that appear in the d block are called the transition elements.
These elements are representative
wide range
physical
andfilling
chemical
Among
They markofthea transition
fromofthe
p orbital
order properties.
to the d orbital
Figure 3.21 The long form of the
them there are highly reactive, moderately reactive and non reactive elements. While most are
periodic
table,atwith
the four
energy
filling
order.
the same
the f block
solids
room
temperature,
roughly
oneByquarter
of reasoning,
them are gases,
one iselements
a liquid. are called
sublevel
blocks identified.
the dinner
transition
elements,
because theyelements.
mark a transition from the
that appear in the
block
are called
the transition
• Elements
d orbital
filling
order
to the f orbital
filling order.
the
inner
transition
elements
• the f block elements are called
18
(VIIIA)
1
(IA)
1 1s
13 14 15 16 17
(IIIA) (IVA) (VA) (VIA) (VIIA)
2
(IIA)
2
2s
3
3s
4
4s
5
5s
6
6s
7
2p
3
4
5
6
7
(IIIB) (IVB) (VB) (VIB) (VIIB)
8
9 10
(VIIIB)
11 12
(IB) (IIB)
3p
3d
4p
4d
5p
4f
5d
6p
7s
5f
6d
s block
(main group
elements)
f block
(inner transition elements)
d block
(transition elements)
p block
(main elements)
3
(IIIB)
Chemical Bonding
A water molecule has a bent shape, carbon dioxide is linear. An ammonia molecule looks like a pyramid,
and sulfur hexafluoride is shaped like an octahedron.
All molecules in nature have a specific shape, which is important to their chemistry.
Each nerve cell in the brain communicates with adjacent nerve cells by releasing molecules called
neurotransmitters from one cell to the next. Enzymes are assisting in the chemical breakdown of food in
our digestive system. The aroma of cologne is the result of odorous molecules migrating to specific sites
in our nasal passages.
Each of these situations depends on the ability of one molecule with a specific
shape to “fit” into a precise location with a corresponding shape (a receptor).
The properties of substances derive from the ways in which particles
bond together
Chemical Bonding
Of the about 120 elements that occur in nature or that have been produced synthetically, only the noble
gases exist naturally as single, uncombined atoms.
In nature, systems of lower energy tend to be favored over systems of higher
energy. In other words, lower-energy systems tend to have greater stability than
higher-energy systems.
Bonded atoms, therefore, tend to have lower energy than single, uncombined
atoms
Defintion
Chemical bonds are electrostatic forces that hold atoms together in
compounds and involve the interaction of valence electrons.
Lewis
Structures
RepresentofAtoms
al Using
bonding
involves
the to
interaction
valence electrons—t
To draw
the Lewisthe
structure
of an atom:
s that
occupy
outermost
principal energy level of an at
e used
Lewis structures in previous studies to indicate th
1.replace its nucleus and inner electrons with its atomic symbol
electrons
atoms.
that to
thetheLewis
2.add dotsofaround
the Recall
atomic symbol
to draw
symbolize
atom’s structu
valence
electrons
(many chemists
place theelectrons
dots startingwith
at the its
top atomic
and continue
place
its nucleus
and inner
sym
adding dots clockwise, at the right, then bottom, then left. then begin again
d dots
the atomic symbol to symbolize the atom’
at thearound
top)
s. Many chemists place the dots starting at the top and c
Drawing a Lewis structure for a molecule lets you see exactly how many
dots
clockwise,
at the
right,
bottom,
electrons
are involved
in each
bond,then
and helps
you tothen
keep left.
track After
of the
of valence
electrons
henumber
first four
dots,
you begin again at the top, as shown be
•
Na
Mg
•
•
•
•
Al
•
•
•
Si
•
• •
•
•
P
•
• •
•
•
S
•
• •
•
•
•
Cl
• •
• •
•
•
•
•
Ar
•
•
• •
his chapter, you will use Lewis structures often to represe
es and the simplest formula unit of an ionic solid. Drawi
ructure for a molecule lets you see exactly how many ele
are involved
in each bond,
and helpsAtoms
you to keep track of th
Using
Lewis Structures
to Represent
valence electrons. In the example below, notice that there ar
to show
the
bonding
pairs
of electrons.
Some chemists use d
There
are two
ways
to show the
bonding
pairs of electrons.
only
•Use dots
Other
chemists
show the bonding pairs as lines between ato
Show the bonding pairs as lines between atoms. In this case dots are
•
reserved
forrepresenting
representing
lone
pair (a non-bonding
pair) of
reserved for
a lone apair
(a non-bonding
pair) of electrons
You will see the second example, with lines for bonding pai
often in this textbook.
• •
•
•
O
• •
• •
C
• •
• •
• •
O
• •
•
•
or
•
•
O
C
• •
O
•
•
(four lone p
4 lone pairs
Bonding and the Properties of Substances
Chemists classify substances according to their bonds and th
attraction that exist between their particles. In the following
you will observe and record data about the properties of five
solid represents a particular type of bonding. Experimental e
Ionic Bonding
Ionic bonding occurs between atoms of elements that have large differences in electronegativity
usually a metal (low electronegativity) and a non-metal (high electronegativity).
The units of ionic compounds such as sodium chloride and magnesium fluoride cannot be
separated easily by direct heating of the crystal salts.
The ions that make up the ionic solid are arranged in a specific array of repeating units.
In solid sodium chloride, for example, the ions are arranged in a rigid lattice
structure. In such systems, the cations and anions are arranged so that
the system has the minimum possible energy
Lattice structure of sodium chloride
non-metals. For example, magnesium in Group 2 (IIA) and fluorine in
Group 17 (VIIA) combine to form the ionic compound magnesium
fluoride, MgF2 . Figure 4.1 shows a repeating unit in the crystal model
of large
magnesium
fluoride.
process that results
in the in
formation
ions
Because of the
differences
in The
electronegativity,
the atoms
an ionicofcompound
usually
be illustrated
withthe
an porbital
or with Lewis structures, as
come from thecan
s block
metals and
blockdiagram
non-metals.
shown in Figures 4.2 and 4.3. Use them as a guide for the Practice
Problems below. Through bonding, the atoms of each element obtain a
Mg
F
valence electron configuration like that of the nearest noble gas. In this
case, the nearest noble gas for both ions is neon. This observation reflects
E.g. magnesium
in Group 2 and fluorine in Group 17 combine to
the octet rule.
Ionic Bonding
F
2s
2p
1s
2s
2p
→
→
→
→
→
→
→
→
→
→
1s
F−
MgF2
2p
→
→
→
→
→
→
→
→
→
→
Mg2+
3s
→
→
→
→
→
→
→
→
→
1s
Figure 4.2
2s
2p
→
→
→
→
→
→
→
→
→
→
Mg
→
→
→
→
→
→
→
→
→
→
→
→
→
→
→
→
→
→
→
→
→
form the ionic compound magnesium fluoride (MgF2).
The figure shows
a repeating unit in the crystal
model of
F
F−
magnesium fluoride.1s 2s
2p
1s 2s
1s
2s
2p
The process that results in the formation of ions can be illustrated with Lewis structures
• •
•
Mg
+
F
•
•
•
•
• •
• •
•
F
• •
•
•
F
•
• •
• •
Practice Problems
1. Write electron configurations for the following:
−
Mg
2+
• •
•
F
• •
•
•
−
Figure 4.3
for MgF2
of elements that have large differences in electronegativity — usually a
metal with a very low electronegativity and a non-metal with a very high
electronegativity. The units of ionic compounds such as sodium chloride
and magnesium
cannot be separated
easily by
heating
of the
Because of the large
differencesfluoride
in electronegativity,
the atoms
indirect
an ionic
compound
usually
crystal salts. The ions that make up the ionic solid are arranged in a specifcome from the s block
metals and the p block non-metals.
ic array of repeating units. In solid sodium chloride, for example, the ions
are arranged in a rigid lattice structure. In such systems, the cations and
Mg
anions are arranged so that the system has the minimum possible energy.
FigureF 4.1
Because of the large differences in electronegativity, the atoms in an
magnesium
compound
usually come
from the
block metals
E.g. magnesium inionic
Group
2 and fluorine
in Group
17s combine
to and the p block
non-metals.
For example,
magnesium
form the ionic compound
magnesium
fluoride
(MgF2in
). Group 2 (IIA) and fluorine in
Group 17 (VIIA) combine to form the ionic compound magnesium
The figure shows
a repeating unit in the crystal model of
fluoride, MgF2 . Figure 4.1 shows a repeating unit in the crystal model
magnesium fluoride.
of magnesium fluoride. The process that results in the formation of ions
can be illustrated with an orbital diagram or with Lewis structures, as
shown in Figures 4.2 and 4.3. Use them as a guide for the Practice
Problems below. Through bonding, the atoms of each element obtain a
valence electron configuration like that of the nearest noble gas. In this
case, the nearest noble gas for both ions is neon. This observation reflects
The process that
results
in the formation of ions can be illustrated with the “box” diagram
the octet
rule.
1s
2p
3s
1s
2s
2p
1s
• •
•
1s
F−
•
2p
2s
2p
→
→
→
→
→
→
→
→
→
→
2s
Mg2+
2s
→
→
→
→
→
→
→
→
→
→
2p
→
→
→
→
→
→
→
→
→
1s
F
2s
→
→
→
→
→
→
→
→
→
→
→
→
1s
Mg
F−
→
→
→
→
→
→
→
→
→
F
→
→
→
→
→
→
→
→
→
→
Ionic Bonding
2s
2p
Figure 4.2
MgF2
Ionic Bonding
Practice problem
•Write electron configurations for the following elements:
a. Li+
b.Ca2+
c.Br−
d.O2−
•Draw Lewis structures for these chemical species
•Draw
orbital diagrams (box) and Lewis structures to show how the following pairs of
elements can combine. In each case, write the chemical formula for the product.
e.Li and S
f.Ca and Cl
g.K and Cl
h.Na and N
Properties of Ionic Solids
In general, ionic solids have the following properties:
•crystalline with smooth, shiny surfaces
•hard but brittle
•non-conductors of electricity and heat
•high melting points
•many ionic solids are also soluble in water (MgF2 is an exception)
The amount of energy given off when an ionic crystal forms from the gaseous ions of its elements is
called the lattice energy (e.g. The lattice energy of MgF2 is 2957 kJ/mol). The same amount of energy
must be added to break the ionic crystal back into its gaseous ions.
tion favours a minimum energy for the system, and constitutes the
−
nt bond between the two hydrogen atoms. Unlike ionic bonding,
ich electrons behave as if they are transferred from one atom to
bonding
results
from the
balance
between the forces of attraction
and
optimum
er, Covalent
covalent bonding
involves
the sharing
of pairs
of electrons.
separation
repulsion
that
act between
the nuclei
and electrons
two or more atoms.
Chapter
3 you
learned
how the quantum
mechanical
model of
applies
atom. The model can also be extended to explain bonding. A
electron4.4 Covalent bonding
Figure
Bonding
nt bond may form when two half-filled atomic orbitals from two
involves forces of attraction
and
H2 molecule
onding involves a balance between the forces of attraction
overlap to share the same region of space. A covalent bond involves
nucleusrepulsion that occur simultaneously
on Example
that act between the nuclei and electrons of two or
rmation
of a new there
orbital,
caused
by the
overlapping
ofhydrogen
atomic
In Hidea
is an
separation
two
−
2 molecule
s. This
is represented
inoptimum
Figure 4.4,
with afor
molecule
of
s. The new
orbitaltheir
hasnucleus-electron
energy levels that
are lower
than those of
nucleusattraction
H2 . atoms
There at
is which
an optimum
separation forattractions,
two hydrogen
atoms
iginal
atomic
orbitals.and
Since
electrons tend
to occupy
the lowest
nucleus
repulsions,
electron-electron
repulsions
achieve
this
repulsion
eir nucleus-electron attractions, nucleus-nucleus repulsions,
+
+
ble balance.
energy level, the new orbitals provide a more energetically
n-electron repulsions achieve this balance. This optimum
This
optimum separation
a minimum
the
able
configuration
than thefavors
two atoms
had energy
before for
they
interacted.
favours a minimum energy for the system, and constitutes the
system, and constitutes the covalent bond between the two
−
nd
between
the
two
hydrogen
atoms.
Unlike
ionic
bonding,
hydrogen of
atoms.
acteristics
Covalent Bonding
ectrons behave as if they are transferred from one atom to
•
ny cases, electron-sharing enables each atom in a covalent bond to
A
optimum
H
H
valent bonding involves the sharing of pairs of electrons.
separation
e a noble gas configuration. For a hydrogen molecule, each atom
pter 3 you learned how the quantum mechanical model applies
es a filled valence level like that of helium by treating the shared
. The model can also be extended to explain bonding. A
Figure 4.4 Covalent bonding
f electrons as if it is part of its own composition. As you can see in
nd may form when two half-filled atomic orbitals from two
involves forcesoverlapping
of attraction
and
region
of increased
of atomic
4.5, a single shared pair of electrons—a bonding pair — fills the
B
repulsion that occur
simultaneously.
lapUnlike
to share thebonding,
same region of space.
A covalent
bond involves
orbitals
electron
density
involves
e level ofionic
both hydrogencovalent
atoms at bond
the same
time. the sharing of pairs of
on of a new orbital,
caused
by the overlapping
of orbital,
atomic caused by the
and the
formation
of fluorine
a newmust
he electrons
period 2 non-metals
from
carbon to
fill their 2s and
e new
orbital
has
energy
levels
that
are
lower
than
those
of
overlapping
oftoatomic
orbitals.
hree
2p orbitals
acquire
a noble gas configuration like that of
atomic orbitals. Since electrons tend to occupy the lowest
Covalent bonding that involves these elements obeys the octet rule.
nergy level, the new orbitals provide a more energetically
formation of the diatomic fluorine molecule, F2 , for example, the
configuration than the two atoms had before they
interacted.
ng (shared) pair of electrons gives each fluorine atom a complete
Covalent Bonding
1s
1s
acquire a noble gas configuration. For a hydrogen molecule, each atom
Characteristics of Covalent Bonding
acquires a filled valence level like that of helium by treating the shared
Generally,
electron-sharing
enables
each
atom
in a covalentAs
bond
acquire
a
pair
of electrons
as if it is part
of its
own
composition.
youtocan
see in
noble gas configuration.
Figure 4.5, a single shared pair of electrons—a bonding pair — fills the
valence level of both hydrogen atoms at the same time.
TheThe
period
2 non-metals
from carbon
to fluorine
must
their 2s and
2p orbitals
to acquire
a
period
2 non-metals
from
carbon
tofillfluorine
must
fill their
2s and
noble gas configuration like that of Ne (octet rule).
their
three 2p orbitals to acquire a noble gas configuration like that of
neon.
bonding
involves
elements
obeys the
octetpairrule.
E.g. InCovalent
the formation
of the that
diatomic
fluorinethese
molecule,
F2, the bonding
(shared)
of
gives eachoffluorine
atom a complete
valence
level. loneF2pairs,
not involved
Inelectrons
the formation
the diatomic
fluorine
molecule,
, forare
example,
thein
bonding.
bonding (shared) pair of electrons gives each fluorine atom a complete
valence level.
bonding pair
lone pairs
• •
•
•
F F
• •
• •
• •
•
•
• •
•
•
or
•
•
F
• •
• •
F
•
•
• •
Each fluorine atom also has three unshared pairs of electrons. These
pairs of electrons, called lone pairs, are not involved in bonding.
The covalent bond that holds molecules of hydrogen, fluorine, and
hydrogen fluoride together is a single bond. It involves a single bonding
ctrons.
Some
molecules
are
bonded
together
withdio
tw
ctrons.
These
are
called
double
bonds.
Carbon
Characteristics of Covalent Bonding
ectrons.
These
are called
double
bonds.
Carbonbonds
dioxi
aSome
covalent
molecule
that
consists
of
double
molecules are bonded together with two shared pairs of electrons. These are called
f adouble
covalent
of double
bonds.
bonds. CO molecule
is an example of a that
covalent consists
molecule that consists
of double bonds
2
• •
O
• •
• • • •
• • • •
• •
O
• •
C
C
• •
• •
O
• •
• •
O
or
• •
• • • •
O
• •
O
or
• •
• •
• •
• •
C
C
• •
O
• •
O
• •
• •
that are bonded with three shared pairs of electron
that are bonded with three shared pairs of electrons
N
,
another
diatomic
molecule,
is
a
tri
s.Molecules
Nitrogen,
2
that are bonded with three shared pairs of electrons have triple bonds. Nitrogen,
ds. Nitrogen, N2 , another diatomic molecule, is a trip
N2, another diatomic molecule, is a triple-bonded molecule
•
•
•
N
N
N N
• •
•• ••
•• ••
• •
•
••
•
or
or
NN NN
•
• •
•
•
• •
•
Chapter4 4Stru
St
Chapter
Characteristics of Covalent Bonding
Bond energy is the energy required to break the force of attraction
between two atoms in a bond and to separate them. Thus, bond energy
Bond energy
is the energy required to break the force of attraction between
is a measure of the strength of a bond. You might expect that the bond
two atoms
in a bond
to separate
Thus,were
bondshared
energybetween
is a measure
of the
energy
wouldand
increase
if morethem.
electrons
two atoms
strength ofbecause
a bond.there would be an increase in charge density between the nuclei
of the bonded atoms. In other words, you might predict that double bonds
The bondareenergy
increases
if more
electrons
shared
twothan
atoms
stronger
than single
bonds,
and that are
triple
bondsbetween
are stronger
because double
there bonds.
is an The
increase
in Table
charge
the nuclei of the
data in
4.1 density
support between
this prediction.
bonded atoms.
s in
alent
Table 4.1 Average Bond Energies of Bonds
Between Carbon Atoms and Between Nitrogen Atoms
Bond
Bond energy (kJ/mol)
Bond
Bond energy (kJ/mol)
C
C
347
N
N
160
C
C
607
N
N
418
C
C
839
N
N
945
Predicting Ionic and Covalent Bonds
There are several methods you can use to predict the type of bond in
an unknown substance. For example, you can consider the substance’s
physical properties. In contrast to ionic solids, covalent (molecular)
Properties of Ionic Solids
In contrast to ionic solids, covalent compounds typically have the following
properties:
• exist as a soft solid, a liquid, or a gas at room temperature
• have low melting points and boiling points
• are poor conductors of electricity, even in solution
• may not be soluble in water
Diamond (C)
Quartz (SiO2)
imply that covalent bonds
are weaker than ionic bonds?
Give evidence to justify your
between
the bonding atoms
answer.
Predicting Covalent or Ionic Bonding
We can use the electronegativity difference
predict the type of bond.
3.3
E.g.
•Two atoms with identical electronegativities, such as chlorine
(∆EN=3.16−3.16=0) share their electrons equally. They are
bonded covalently.
•In sodium chloride, chlorine (EN=3.16) attracts an electron
much more strongly than sodium (EN=0.93). Therefore, sodium’s
valence electron has a very high probability of being found near
chlorine. A high electronegativity difference is characteristic of
ionic compounds.
For atoms that have ∆EN between 0.4 and 1.7, the bond is polar
covalent. A polar covalent bond has an unequally shared pair of
electrons between two atoms.
This unequal sharing results in a bond that has partially positive
and partially negative poles.
mostly ionic
(∆EN > 1.7)
+
−
2.0 ∆EN
polar covalent
(∆EN 0.4 –1.7)
mostly covalent
(∆EN < 0.4)
0
The
relationship
bonding character
Figure
4.6 between
The
relationship
and electronegativity difference
between bonding character and
to
Metallic Bonding
About two-thirds of all the naturally occurring elements are metals (see
lesson 1).
Metals conduct electricity and heat in both their solid and liquid states.
Most metals are malleable and ductile (can be easily stretched, bent, and
deformed without shattering of the whole solid).
In general, metals change state at moderate to high temperatures. Most metals
have either one or two valence electrons
Sodium
Lithium
Potassium
Metallic Bonding
Based on electronegativity differences metals do not form ionic bonds with other metals.
Similarly, metals do not have a sufficient number of valence electrons to form covalent bonds
with one another.
Metals do, however, share electrons. Unlike the electron sharing in covalent compounds,
however, electron sharing in metals occurs throughout the entire structure of the metal.
Metals are composed of a densely packed core of metallic
cations, within a delocalized region of shared, mobile
valence electrons (free-electron model).
The force of attraction between the positively charged cations
and the pool of valence electrons that moves among them
constitutes a metallic bond.
Co
Na
millions
of atoms
e− sea
M
A
Properties of the Metallic Bonding
The free-electron model explains many properties of metals:
•Conductivity:
Metals are good conductors of electricity and heat because electrons can move freely
throughout the metallic structure. This freedom of movement is not possible in solid ionic
compounds, because the valence electrons are held within the individual ionic bonds in the
lattice.
•Melting and Boiling Points:
The melting and boiling points of Group 1 metals are generally lower than the melting and
boiling points of Group 2 metals. Because the greater number of valence electrons and the
larger positive charge of Group 2 atoms result in stronger metallic bonding forces
Despite this great diversity, metals haveBased
man
Figure 4.8 A representation of
Metals conduct electricity and heat to
in form
both io
th
the free-electron model of metallic
Most
is
The free-electron model explains many properties
of metals
metals: are malleable and ductile;
not.that
Simil
bonding. This model applies to
bent, and deformed without shattering
of the
to form
co
metal alloys as well as to metallic
+ + + of heat
and Ductility:
•Malleability
made to be with the+addition
or press
electrons.
elements.
+ +and
+ boiling
+ The points,
The malleability of metals can be explained as metallic
bondsofare
non-directional.
positive ins
broad range
melting
electron
external
ions are layered as fixed arrays (like soldiers linedatup
for inspection).
to
+temperatures.
+ stress
+ + is applied
moderate
to highWhen
Most
meta
metal.
Th
stress applied
a metal, one layer of positive ions can slide over
another
layer.
The
layers
move
without
+ + kind
+ + of bonding mod
valence electrons. What
being com
breaking the array.
+ + + +
properties?
delocalize
Based on electronegativity differences,
yo
attraction
to form ionic bonds with themselves
or with
electrons
+
not. Similarly, metals do not have a sufficien
+ +
Properti
to form covalent bonds with one another.
Me
+
+
+
deformed
metal
+ + + +
electrons. Unlike the electron
sharing
cova
Theinfree-e
+
+
+
+
+ + + +
electron sharing in metals
occurs
throughout
external
+
+
+
+
• Conduc
+ + + +
metal. The free-electron
model
shown in Fig
stress applied
+
+
+
electron
+ + + +
being composed of a+ densely
packed
core of
+
of
move
+ + + +
delocalized region of
shared,
mobile
valence
+
valence
attraction between the positively charged cat
• Malleab
electrons
that
moves
among
them
constitutes
Figure 4.9 Metals are easily
+
by view
deformed because one layer of pos+ +
are ofte
Properties
Explained
by the
Free-Electr
itive
ions
can
slide
over
another.
At
+
+
+
deformed metal
When s
the
same
time,
the
free
electrons
The
free-electron
model
explains
many
prop
+ + + +
over an
(shown as a yellow cloud) continue
Properties of the Metallic Bonding