Chemical Bonding
What is a Chemical Bond?



a bond is a force of attraction
between two atoms
the source of the force are attractions
and repulsions between electrons and
the atomic nucleii
before we get into detail, do the Ball &
Stick activity, to get an idea of what
molecules look like.
Review of Grade 11




Quantum theory
electron configuration
orbital diagrams
Lewis diagrams
Quantum Mechanics and
Atomic Orbitals

Quantum mechanics is
a mathematical
treatment into which
both the wave and
particle nature of matter
could be incorporated.
Quantum Mechanics


since the electron is both a
wave and a particle it is
impossible to give it’s
location or speed with
certainty.
gives a probability density
map of where an electron
has a certain statistical
likelihood of being at any
given instant in time.
Quantum Numbers


The probability map reveals the atomic
orbitals, and their corresponding
energies.
An orbital is described by a set of three
quantum numbers.
Principal Quantum Number, n



This relates to the energy of the
electron
As n becomes larger, the atom
becomes larger and the electron is
further from the nucleus.
This is directly related to the period of
the atom on the Periodic Table
Angular momentum quantum
number, l


This quantum number defines the shape
of the orbital.
There are 4 shapes:





s
p
d
f
-
begins at n = 1
begins at n = 2
begins at n = 3
begins at n = 4
Theoretical g, h, i, etc. orbitals exist, but
no atoms have been created to use them.
Magnetic Quantum Number,
ml

Magnetic quantum numbers give the
three-dimensional orientation of
each orbital.




s
p
d
f
-
has 1 orientation
has 3 orientations
has 5 orientations
has 7 orientations
s Orbitals

Spherical in shape.
p Orbitals


propeller shaped
Have two lobes with a node between them.
d Orbitals

Four of the
five orbitals
have 4 lobes;
the other
resembles a p
orbital with a
doughnut
around the
center.
f orbitals (flowers)
Spin Quantum Number, ms



electrons have spin,
which creates a
magnetic field
there are two spin states
possible, +1/2 and -1/2
a single orbital can
hold a maximum of two
electrons, which must
have opposite spin (Pauli
Exclusion Principle).
Electron Address

thus every electron location is defined
in terms of 4 things:
a)
b)
c)
d)
Principal Quantum Number - 1 to 7
Angular Quantum Number – s, p, d or f
Magnetic Quantum Number – implied by
number of electrons in each shape;
s has 2, p has 6, d has 10 and f has 14
Spin Quantum Number – why each
orbital can contain 2 electrons
Electron Configurations


Electrons tend to occupy the lowest
energy orbitals.
Thus the electron configuration of an atom
is the arrangement of the electrons from
the lowest energy level to the highest.
Electron Configurations



Consist of
 Number denoting the energy level.
 Letter denoting the type of orbital.
 Superscript denoting the number of electrons
in those orbitals.
For instance:
 Iron (Fe) – contains 26 electrons
 1s22s22p63s23p64s23d6
watch the order of filling
Electron Configuration




Potassium - 19 electrons
 1s22s22p63s23p64s1
Silver - 47 electrons
 1s22s22p63s23p64s23d104p65s24d9
Tungsten - 74 electrons
 1s22s22p63s23p64s23d104p65s24d105p66s24f145d
4
Plutonium - 94 electrons
 1s22s22p63s23p64s23d104p65s24d105p66s24f145d
10
6p67s25f6

Write the correct electron configuration for
the following:
 Si, S, P, Ca, As, Fe, Br, Kr, At, U, Na1+,
F1-, Ne
Electron Configuration













Si - 14 e1- 1s22s22p63s13p3
S - 16 e1- 1s22s22p63s23p4
P - 15 e1- 1s22s22p63s23p3
Ca - 20 e1- 1s22s22p63s23p64s14p1
As - 33 e1- 1s22s22p63s23p64s23d104p3
Fe - 26 e1- 1s22s22p63s23p64s23d6
Br - 35 e1- 1s22s22p63s23p64s23d104p5
Kr - 36 e1- 1s22s22p63s23p64s23d104p6
At - 85 e11s22s22p63s23p64s23d104p65s24d105p66s24f145d106p5
U - 92 e11s22s22p63s23p64s23d104p65s24d105p66s24f145d106p67s
25f4
Na1+ - 10 e11s22s22p6
F1- - 10 e1- 1s22s22p6
Ne - 10 e1- 1s22s22p6
Electron Promotion





promotion of an outer ‘s’ electron to the
adjacent ‘p’ orbital.
turns non-bonding electrons into bonding
electrons
allows atoms to make more chemical bonds
and achieve a lower energy
applies to elements from groups 2, 13 and 14
only
for these elements promotion is the rule
Electron Promotion
Element
Unhybridized
Hybridized
beryllium
1s22s2
1s22s12p1
boron
1s22s22p1
1s22s12p2
carbon
1s22s22p2
1s22s12p3
Orbital Diagrams


are another way to illustrate the position of
electrons.
They are best learned by comparison with electron
configuration:

Na (11 protons, 11 electrons)
electron configuration:
1s22s22p63s1

orbital diagram:

1s
2s
2p
3s
↑↓
↑↓
↑↓ ↑↓ ↑↓
↑
Orbital Diagrams
Group
Representative
Element
Electron configuration
Orbital Diagram
1
lithium
1s22s1
1s
↑↓
2s
↑
2p
2
beryllium
1s22s12p1
↑↓
↑
↑
13
boron
1s22s12p2
↑↓
↑
↑ ↑
14
carbon
1s22s12p3
↑↓
↑
↑ ↑ ↑
15
nitrogen
1s22s22p3
↑↓
↑↓
↑ ↑ ↑
16
oxygen
1s22s22p4
↑↓
↑↓
↑↓ ↑ ↑
17
fluorine
1s22s22p5
↑↓
↑↓
↑↓ ↑↓ ↑
18
neon
1s22s22p6
↑↓
↑↓
↑↓ ↑↓ ↑↓
Repeat the last assignment, giving the orbital diagrams for the elements.
• Na1+, F1-, Ne are all the same:
Electron dot (Lewis)
diagrams

gives information only concerning the valence
electrons.






the electrons on the outside of an atom
the electrons responsible for bonding
the electrons gained or lost when an atom ionizes.
the electrons in the s and p orbitals of the highest
energy level reached by the electrons of an atom.
In this class when valence electrons are
mentioned, the only elements concerned are
those in groups 1, 2, and 13 through 18.
all atoms in the same group have the same
Lewis diagram !!
Lewis diagrams

group 1

group 2

group 13

group 14

group 15

group 16

group 17

group 18
Lewis diagrams

Repeat last assignment, making lewis
diagrams.
Lewis Diagram Answers:
Fe and U have no lewis diagrams
What do Orbital and Lewis
diagrams tell us?

both give information about valence
electrons.


if valence electrons are paired, they
cannot be used for bonding with other
atoms. They are lone-pair electrons.
unpaired valence electrons are bonding
electrons.
• 5 valence electrons
• 1 lone pair
• 3 bonding electrons; this atom makes 3
chemical bonds.
• 4 valence electrons
• 0 lone pair
• 4 bonding electrons; this atom makes 4
chemical bonds
• 7 valence electrons
• 3 lone pair
• 1 bonding electrons; this atom makes 1
chemical bond
• 8 valence electrons
• 4 lone pair
• 0 bonding electrons; this atom makes 0
chemical bonds
Valence level expansion

Some compounds occur which cannot be
easily explained:
PF5, SF6, ClF7, ArF8


in each case the number of chemical bonds
is equal to the number of valence electrons.
this can only happen if electrons are
promoted to a higher energy level. In this
case it is the adjacent ‘d’ orbital.
PF5

the normal orbital diagram looks like
this:

with valence level expansion it looks
like this:
Valence level expansion

includes:



elements of groups 15 to 18, from period
3 down
periods 1 and 2 do not have a ‘d’ orbital
to promote to.
Please note that valence level
expansion is the exception, not the
rule.
Chemical Bonding





is all about the electrons.
in most of our discussions we will
concentrate the valence electrons.
since there are 4 of these orbitals in any
quantum it requires 8 electrons (an octet) to
fill them.
the tendency of atoms to try to fill out the
outer ‘s’ and ‘p’ shells is the octet rule.
exception: atoms from groups 2 and 13
Chemical Bonding

Bonding can happen in one of two ways:

To share electrons 

The outer orbitals of 2 atoms overlap so that
each atom is in the vicinity of a full set of
valence electrons.
This type of bonding is called covalent
bonding.
Chemical bonding




Atoms gain or lose electrons to arrive at a
full set of valence electrons.
When atoms gain or lose electrons they
become ions.
Ions are attracted to ions of opposite charge
and repelled by ions of the same charge.
This type of bonding is called ionic bonding
How Many Bonds ???
• the number of bonds made by atoms in
the ‘s’ and ‘p’ blocks of the Periodic Table
is determined by how many electrons they
are away from an octet:
How Many Bonds ???
Group
Valence
Electrons
Covalent
Bonds
Ionic
Bonds
1
2
13
1
2
3
1
21
31
1 (1+)
2 (2+)
3 (3+)
14
15
16
17
18
4
5
6
7
8
41
3 (52)
2 (62)
1 (72)
0 (82)
3 (3-)
2 (2-)
1 (1-)
0
1electron
promotion
2valence
level expansion
• most of our discussion will be centred on
covalent bonding.
Electronegativity



is a measure of how strongly an atom is
holding on to its valence electrons.
If an atom loses an electron fairly easily it
has a low electronegativity (and tends to be
a cation).
If an atom tends not to lose electrons, but
tends to steal them from other atoms (and
become an anion) it has a high
electronegativity.

To determine what type of bond exists
between two atoms you subtract their
respective electronegativities:


if the electronegativity difference is 0.2
or less, the bond is covalent
if the electronegativity difference is 1.7
or greater the bond is ionic.
Electronegativity


If the electronegativity difference between two
atoms is between 0.3 and 1.6 the bond is polarcovalent.
The greater the electronegativity difference the
greater the ionic character of the bond:
Polar-Covalent Bonding


Covalent bonding implies equal
sharing of electrons.
If sharing is not equal, the electrons in
a bond will spend more time with one
atom than the other.
Polar-Covalent Bonding


The atom where the electrons spend
more time will have a net negative
charge, while the atom at the other
end of the bond will be positive.
This type of bond is polar-covalent.
Assignment

Determine the electronegativity difference for each
chemical bond. If the bond is polar covalent draw an
arrow in the direction of the dipole, from positive to
negative:
C-H
END = | 2.5 - 2.1 | = 0.4
polar covalent bond
N-H
P-H
Br - Cl
B-F
Si - Cl
O-H
S-O
Cu - Br
C - Cl
N-I
C-O
N-H
• END = |3.0 – 2.1| = 0.9
• polar covalent bond:
N-H
B-F
• END = |2.0 – 4.0| = 2.0
• ionic bond:
[B]3+[F]1S–O
• END = |2.5 – 3.5| = 1.0
• polar covalent bond:
S-O
P-H
• END = |2.1 – 2.1| = 0.0
• covalent bond:
P–H
Si - Cl
• END = |1.8 – 3.0| = 1.2
• polar covalent bond:
Si - Cl
Cu - Br
• END = |1.9 – 2.8| = 0.9
• polar covalent bond:
Cu - Br
N–I
• END = |3.0 – 2.5| = 0.5
• polar covalent bond:
N-I
Br – Cl
• END = |2.8 – 3.0| = 0.2
• covalent bond:
Br – Cl
O–H
• END = |3.5 – 2.1| = 1.4
• polar covalent bond:
O–H
C – Cl
• END = |2.5 – 3.0| = 0.5
• polar covalent bond:
C – Cl
C–O
• END = |2.5 – 3.5| = 1.0
• polar covalent bond:
C–O
Back to the Ball & Stick

determine the electronegativity
difference and bond type for each
bond in each of the molecules in the
activity.
Covalently-Bonded Structures



we now have to consider molecules
made of several atoms.
most of the following discussion will
concern itself with molecules made
with covalent or polar-covalent bonds.
ionic bonds (and others) will return
later in the unit.
Lewis Structures


Lewis structures allow us to predict
how atoms will come together to make
molecules.
Lewis structures are representations of
molecules showing all electrons,
bonding and nonbonding.
Writing Lewis Structures
1.
PCl3
Find the sum of
valence electrons of all
atoms in the
polyatomic ion or
molecule.

5 + 3(7) = 26

If it is an anion, add one
electron for each
negative charge.
If it is a cation, subtract
one electron for each
positive charge.
Writing Lewis Structures
2.
Keep track of the electrons:
26  6 = 20
The central atom is
the least
electronegative
element that isn’t
hydrogen. Connect
the outer atoms to it
by single bonds.
Writing Lewis Structures
3.
Keep track of the electrons:
26  6 = 20  18 = 2
Fill the octets of the
outer atoms.
Writing Lewis Structures
4.
Keep track of the electrons:
26  6 = 20  18 = 2  2 = 0
Fill the octet of the
central atom.
Writing Lewis Structures
5.
If you run out of
electrons before the
central atom has an
octet…
…form multiple bonds
until it does.
Exceptions to Octet Rule

Electron Promotion



group 2 central atom will have 4 electrons
group 13 central atom will have 6 electrons
Valence Level Expansion




group 15 central atom will have 10 electrons
group 16 central atom will have 12 electrons
group 17 central atom will have 14 electrons
group 18 central atom will have 16 electrons
Writing Lewis Structures

Write Lewis Structures for the following
molecules:
F2
BCl3
SeH2
MgH2
SiF4
C 3H 8
C 3H 6
C3H4
Molecular Shapes


The shape of a
molecule plays an
important role in its
reactivity.
By noting the number
of bonding and
nonbonding electron
pairs we can easily
predict the shape of
the molecule.
What Determines the Shape of
a Molecule?


Simply put, electron pairs, whether they be
bonding or nonbonding, repel each other.
By assuming the electron pairs are placed as far
as possible from each other, we can predict the
shape of the molecule.
Valence Shell Electron Pair
Repulsion Theory (VSEPR)
“The best
arrangement of a
given number of
electron domains is
the one that
minimizes the
repulsions among
them.”
Molecular Geometries

the shape of any molecule can be
predicted based on two things:
1.
2.
Group number of the central atom
number of atoms bonded to the central
atom
Group 2 Geometries
Central
Atom
Magnesium


Bonding Lone Pair Bond Type Shape Example
Electrons Electrons
2
0
all single
linear
MgI2
one double
linear
MgO
In this domain, there is only one molecular
geometry: linear.
NOTE: If there are only two atoms in the
molecule, the molecule will be linear no
matter what the electron domain is.
Group 13 Geometries
Central Bonding Lone Pair
Atom Electrons Electrons
Boron

3
0
Bond
Type
Shape
Example
all single
trigonal
planar
BI3
one
double
linear
BIO
one triple
linear
BN
because there are no lone pair electrons the
molecular is planar (flat).
Group 14 Geometries
Central Bonding Lone Pair
Atom Electrons Electrons
Carbon
4
0
Bond
Type
Shape
Example
all single
tetrahedral CH4
one
double
trigonal
planar
one triple, linear
or
two
double
COH2
HCN
CO2
Summary for Groups 2, 13 & 14

because there are no lone pair
electrons:




central atom bonded to 1 atom: linear
central atom bonded to 2 atoms: linear
central atom bonded to 3 atoms:
trigonal planar
central atom bonded to 4 atoms:
tetrahedral
Group 15 Geometries
Central
Atom
Nitrogen
Bonding Lone Pair
Electrons Electrons
3
1
Bond
Type
Shape
all single
trigonal
pyramidal
NH3
one
double
angular
NOH
one triple linear
5
0
Example
all single
N2
trigonal
NCl5
bipyramidal
Group 16 Geometries
Central
Atom
Oxygen
Bonding Lone Pair
Electrons Electrons
2
6
2
0
Bond
Type
Shape
Example
all single
angular
H2O
one
double
linear
O2
all single
octahedral OF6
Group 17 & 18 Geometries

group 17



normally makes 1 chemical bond (linear)
with valence level expansion can make 7
bonds (ClF7).
group 18


normally makes no bonds
with valence level expansion can make 8
bonds (ArF8).
Larger Molecules


In larger molecules, it makes more sense to talk
about the geometry about a particular atom rather
than the geometry of the molecule as a whole.
This molecule is tetrahedral about the first carbon,
planar trigonal about the second and angular about
the single-bonded oxygen.
Larger Molecules
This approach
makes sense,
especially because
larger molecules
tend to react at a
particular site in the
molecule.
Using VSEPR Theory

Determine the shape of each
molecule, or atom within the molecule:
F2
PCl3
SeH2
MgH2
SiF4
C 3H 8
C 3H 6
C3H4
Polarity


previously we
discussed bond
dipoles.
But just because a
molecule possesses
polar bonds does not
mean the molecule as
a whole will be polar.
Polarity


By adding the
individual bond
dipoles, one can
determine the overall
dipole moment for
the molecule.
In other words we can
see if the dipoles
reinforce each other,
or cancel out.
Polarity
Assignment

Perform the following activities for the
molecules in the Balls & Sticks Activity:
1.
2.
3.
4.
Draw a lewis structure
Draw a structural diagram and indicate shape.
Determine the electronegativity difference for
each bond type and determine whether the
bond is covalent, polar covalent or ionic
Determine if they will be polar, or non-polar.
Properties of Molecules
Intermolecular Forces
Properties of Molecules



are determined by how the molecules
interact with each other.
How they interact is determined by the
forces of attraction between molecules.
These are called intermolecular forces;
the forces which act between molecules, to
draw them together, forming the various
phases of matter.
How do we measure force ?




the best way is by temperature.
temperature is a measure of kinetic energy;
the lowest temperature represents zero
kinetic energy.
for a substance to melt or boil the kinetic
energy must overcome the intermolecular
force.
the higher the melting temperature or boiling
temperature, the greater the intermolecular
force.
Van der Waals Forces


occur between covalently bonded
molecules.
three kinds:



London disersion
dipole-dipole attraction
hydrogen bonding
London Dispersion Forces



are the dominant forces between
covalently bonded, non-polar
molecules
based on the formation of
instantaneous dipoles.
the more electrons in a molecule, the
stronger the force.
Relationship of Boiling Point to Number of
Electrons and Molar Mass
Melting Point Boiling Point
# e°C
°C
CH4
C2H6
C3H8
C4H10
C8H18
10
18
26
34
66
- 182
- 183
- 190
- 138
- 57
-161
- 88
- 44
- 0.5
+125
London dispersion forces

are influenced by
shape:

Normal Pentane
(C5H12)



m.p. -130C,
b.p. 36°C
Neopentane (C5H12)


m.p. -20°C,
b.p. 9°C
Dipole-Dipole attraction



is a force which acts between polar
molecules (ex. H2S).
results from the attraction of the opposite
poles of the permanent molecular dipoles.
These substances generally have higher
melting and boiling points than non-polar
molecules with similar molecular weights (or
numbers of electrons).
Ion-Dipole Interactions


A fourth type of force, ion-dipole interactions
are an important force in solutions of ions.
The strength of these forces are what make
it possible for ionic substances to dissolve in
polar solvents.
Dipole-Dipole Interactions
The more polar the molecule, the higher
is its boiling point.
How Do We Explain This?


The nonpolar series
(SnH4 to CH4)
follow the expected
trend.
The polar series
follows the trend
from H2Te through
H2S, but water is
quite an anomaly.
Hydrogen Bonding


a specialized form of dipole-dipole attraction.
It occurs as when O, N, and F are bonded to
H, owing to the large electronegativity
difference:
 O - H
3.5 - 2.1 = 1.4
 N - H
3.1 - 2.1 = 1.0
 F - H
4.1 - 2.1 = 2.0
Hydrogen Bonding



This is a stronger force than standard
dipole-dipole attraction.
Molecules with hydrogen bonding will
have boiling points and melting points
quite a bit higher than molecules that
have only dipole-dipole or London
dispersion forces.
Hydrogen bonding is responsible for
many of the unusual properties of
water.
Hydrogen bonding


is responsible for folding, final
structure and function of proteins
holds the DNA strands together, but
allows them to be unzipped for copying
and gene expression.
Relationship Between Polarity, Force
and Boiling Point
# e- B.P. (°C) Polarity
C2H6 18
- 161
H2S 18
- 60
H2O 10
+ 100
Force
Non-polar London
Dispersion
Polar
Dipole
- Dipole
Very polar Hydrogen Bonds
In Summary


Non-Polar Molecule
 London Dispersion Force
 More electrons, higher B.P., M.P.
 If isoelectronic, more compact higher M.P., lower B.P.
Polar Molecule
 Dipole-Dipole Attraction
 Higher B.P. & M.P. than London
 If both polar, higher electronegativity difference has
higher B.P. & M.P
 If have O-H, N-H, or F-H bond, hydrogen bonding,
much higher M.P. & B.P
Ionic Bonding




generally occurs in compounds of metals
and non-metals (salts).
It is the result of the attraction of oppositely
charged ions.
The structures formed are very orderly and
are given the name crystal lattice.
Ionic solids are called crystals.
Ionic Bonding




No sharing of electrons occurs
between the ions in the crystal lattice.
As a result, ionic solids are brittle.
Ionic solids conduct electricity only in
the molten state, and not very well.
Ionic solids are characterized by very
high melting and boiling points.
Metallic Bonding


is the bonding which occurs between metals
in the Periodic Table.
It is characterized by close packing of the
atoms, with the electrons delocalized; that
is, they are free to jump from atom to atom,
filling unoccupied orbitals.
Metallic Bonding




This free sharing of electrons allows metals
to conduct electricity freely (copper
conducts electricity 100 000 times better
than molten NaCl).
The free electrons also act as a lubricant,
allowing metal atoms to slide over one
another without affecting the integrity of the
material. Thus metals are malleable and
ductile.
This bond is strong, giving most metals high
melting and boiling points.
This bond is also variable.
Network Covalent Bonding







regular covalent bond, expanded to 2 or 3 dimensions in a
network which is theoretically infinite
Network solids include diamonds, graphite, quartz, and
most rocks.
very hard (diamonds are the hardest substance known).
brittle,
do not conduct electricity.
crystals.
very high melting and boiling points
Intermolecular Forces Affect
Many Physical Properties
The strength of the
attractions between
particles can greatly
affect the properties
of a substance or
solution.
Viscosity
• Resistance of a liquid
to flow is called
viscosity.
• It is related to the ease
with which molecules
can move past each
other.
• Viscosity increases
with stronger
intermolecular forces
and decreases with
higher temperature.
Surface Tension
Surface tension
results from the net
inward force
experienced by the
molecules on the
surface of a liquid.
Vapor Pressure
• At any temperature, some molecules in a
liquid have enough energy to escape.
• As the temperature rises, the fraction of
molecules that have enough energy to
escape increases.
Intermolecular Forces in
Summary
Strength of Attraction
Molecular
(London dispersion)
Molecular
(dipole-dipole)
Molecular
(hydrogen bonding)
Ionic
1 to 50
Metallic
20 to 80
Network Covalent
5
15
50
100
Melting and Boiling Point
Molecular
(London dispersion)
very low (nitrogen boils at
- 196 C) VARIABLE
Molecular
(dipole-dipole)
Molecular
(hydrogen bonding)
low (H2S boils at -61 C)
Ionic
high
(NaCl boils at +1413 C)
variable (Hg @ +357 C,
W @ +5660 C)
Metallic
Network Covalent
medium
(H2O boils at +100 C)
very high
(SiO2 boils at 2600 C)
Properties of Solids
Molecular
(London dispersion)
soft, waxy
Molecular
(dipole-dipole)
Molecular
(hydrogen bonding)
more rigid
Ionic
long-range crystalline,
hard, brittle
short-range crystalline,
ductile, malleable
Metallic
Network Covalent
crystalline, brittle
long-range crystalline,
hard, brittle
Conductance of Heat and
Electricity
Molecular
(London dispersion)
non-conductor
Molecular
(dipole-dipole)
Molecular
(hydrogen bonding)
non-conductor
Ionic
conductor in liquid or
dissolved phase
conductor in solid or liquid
phase
Metallic
Network Covalent
non-conductor
non-conductor in 3d form,
some conductance in 2d
Solubility in H2O
Molecular
(London dispersion)
Not soluble
Molecular
(dipole-dipole)
Molecular
(hydrogen bonding)
Soluble
Ionic
Soluble
Metallic
Not soluble
Network Covalent
Not soluble
Soluble
Other Stuff
Coordinate Covalent bonds

is defined as a covalent bond where one
atom provides both of the electrons:
Polyatomic Ions

we can use coordinate covalent bond
theory to explain most ions:
Resonance
This is the Lewis
structure we
would draw for
ozone, O3.
+
-
Resonance

But this is at odds
with the true,
observed structure
of ozone, in
which…


…both O—O bonds
are the same
length.
…both outer
oxygens have a
charge of 1/2.
Resonance


One Lewis structure
cannot accurately
depict a molecule
such as ozone.
We use multiple
structures, resonance
structures, to
describe the
molecule.
Resonance
Just as green is a synthesis
of blue and yellow…
…ozone is a synthesis of
these two resonance
structures.
Resonance


In truth, the electrons that form the second C—O
bond in the double bonds below do not always sit
between that C and that O, but rather can move
among the two oxygens and the carbon.
They are not localized, but rather are delocalized.
Resonance


The organic compound
benzene, C6H6, has
two resonance
structures.
It is commonly
depicted as a hexagon
with a circle inside to
signify the delocalized
electrons in the ring.