Chapter 6 – Chemical Bonding
6-1 Introduction to Chemical
Bonding





Remember, the electrons in the outermost
energy level, the valence electrons are
involved in bonding.
Atoms like to have a filled outer energy
level (valence shell).
Atoms can gain electrons to fill their
valence shell.
Atoms can lose electrons to reveal a full
valence shell.
Atoms can share electrons to get a filled
valence shell.
6-1 Chemical Bond


A mutual electrical attraction
between the nuclei and valence
electrons of different atoms that
binds the atoms together
Most substances on earth are
compounds – all compounds contain
chemical bonds – they hold the
atoms together
6-1 Why do chemical bonds form?



Most atoms by themselves have high
potential energy
They form bonds to minimize their
potential energy
Most atoms are less stable by
themselves than combined with
other atoms
6-1 Types of Chemical Bonds



Ionic bonding –
electrical attraction
between a large
number of cations and
anions – formation of
ions involves transfer
of electrons
Covalent bonding –
sharing of electron
pairs between atoms
to form molecules
Metallic bonding –
delocalized electrons
give metals their
special properties
6-1 Ionic or Covalent?



Bonding between two atoms is not
usually purely ionic or purely
covalent
Remember, electronegativity is a
measure of an atom’s ability to
attract electrons toward itself in a
chemical bond
Percent ionic/covalent character can
be calculated using difference in
electronegativity between two atoms
6-1 Ionic or Covalent?



Find the difference in
electronegativity
between the two
elements in the bond
The difference
(absolute value)
corresponds to the
percentage ionic
character
Chart on right is in
textbook – p. 162
6-1 Ionic or Covalent?




>50% - ionic
50% or less –
covalent
Bonding between two
identical atoms is
completely covalent –
there is NO
electronegativity
difference between
atoms of the same
element – the
electrons are shared
evenly
Called NONPOLAR
COVALENT BOND
6-1 The Diatomic Molecules


Diatomic molecule – a molecule made of two
atoms of the same element
Seven elements exist naturally as diatomic
molecules: H, N, O, F, Cl, Br, I
6-1 Polar Covalent Bonds



Most covalent bonds are not purely covalent
Most covalent bonds are polar covalent – the
electron density is shifted toward the more
electronegative atom (uneven distribution of
charge)
Bonds with between 5 and 50% ionic character
are POLAR COVALENT
6-1 Polar Covalent, Nonpolar
Covalent or Ionic?






H-S
Cs-S
Cl-S
Ca-Cl
O-Cl
Br-Cl
element
electronegativity
H
2.1
S
2.5
Cs
0.7
Cl
3.0
Ca
1.0
O
3.5
Br
2.8
6-2 Covalent Bonding and
Molecular Compounds



Molecule – a neutral group of atoms
that are held together by covalent
bonds
Molecular formula – shows the types
and numbers of atoms combined in a
single molecule of a molecular
compound
BF3, CH3OH, CCl4
6-2 Formation of a Covalent Bond
1.
2.
3.
4.
Atoms far apart,
don’t influence
each other
Atoms approach
each other,
charged particles
begin to interact
(attractions and
repulsions)
Attractive force
dominates until
point 3, when
attraction equals
repulsion
If atoms approach
further, repulsion
becomes
increasingly
greater, PE
increases sharply
6-2 Attractions and Repulsions
Between Atoms
6-2 Characteristics of the Covalent
Bond



BOND LENGTH – distance between
two bonded atoms at their minimum
potential energy
BOND LENGTH REPRESENTS A
POTENTIAL ENERGY WELL!
Forming a bond ALWAYS releases
energy – the same amount of energy
must be ADDED to break the bond –
called BOND ENERGY
6-2 Hydrogen Atoms in H2 Have
Noble Gas Configuration
6-2 The Octet Rule



Chemical compounds tend to form so
that each atom, by gaining, losing or
sharing electrons, has an octet of
electrons in its highest occupied
energy level
F2, HCl, CH4
Most main group elements form
covalent bonds according to the octet
rule
6-2 Exceptions to the Octet Rule

Incomplete octet – BF3

Expanded octet – SF6
6-2 Electron Dot Diagrams

Electron configuration notation in
which only the valence electrons are
shown, indicated by dots placed
around the element’s symbol
6-2 Lewis Structures


Formulas in which atomic symbols
represent nuclei and core electrons,
dot-pairs or lines represent electron
pairs (unshared electron pairs,
bonding electron pairs)
H2, F2, HF, NH3, CH4, CH3I, SH2
6-2 Multiple Covalent Bonds



Two atoms can share one, two or
three pairs of electrons between
them
Double bond – two pairs of electrons
shared between atoms (O2)
Triple bond – three pairs of electrons
shared between atoms (N2)
6-2 Multiple Covalent Bonds


Double bonds generally have
HIGHER bond energies and SHORTER
bond lengths than single bonds
Triple bonds generally have HIGHER
bond energies and SHORTER bond
lengths than double bonds
6-2 Sample Problems

C2H6

C2H4

C2H2
6-2 Resonance Structures


Some molecules or ions cannot be
adequately represented by one Lewis
structure
Resonance structures for ozone, O3
6-2 Covalent-Network Bonding



Do not
contain
individual
molecules
Continuous,
three
dimensional
networks of
bonded
atoms
Ex. graphite
Diamond: A Covalent Network
6-3 Ionic Bonding and Ionic
Compounds


Ionic compound –
composed of positive and
negative ions that are
combined so that the
numbers of positive and
negative charges are equal
Formula unit – simplest
collection of atoms from
which an ionic compound’s
formula can be established
– simplest whole number
ratio of cations to anions
that will give a neutral
formula

calcium fluoride

sodium oxide

magnesium sulfide
6-3 Formation of Ionic Bonds
(represented with dot diagrams)

sodium chloride

calcium fluoride

magnesium sulfide
6-3 Characteristics of Ionic Bonding




Ions arrange
themselves to
minimize potential
energy
Oppositely charged
ions attract each
other
Cations surrounded
by anions and vice
versa
Arrangement is
called a crystal
lattice
6-3 Strength of Ionic Bonds





Bond formation
releases energy
Lattice energy –
energy released when
one mole of an ionic
crystalline compound
is formed from
gaseous ions
Negative values
indicate energy is
released
More negative =
stronger bond
Table 6-3 (p. 179)
compound
Lattice energy
(kJ/mol)
NaCl
NaBr
CaF2
CaO
LiCl
LiF
MgO
KCl
-787.5
-751.4
-2634.7
-3385
-861.3
-1032
-3760
-715
6-3 Ionic vs. Molecular Compounds


Forces that hold
ions together are
very strong
Covalent bonds
also very strong,
but forces of
attraction between
molecules
(intermolecular
forces) much
weaker
6-3 Ionic v. Molecular Compounds
Ionic
Forces
Melting Point
Boiling Point
Hardness
Molecular
6-3 Why are ionic compounds
brittle?

Shifting ions slightly puts like charges next
to each other
6-3 Solubility of Ionic Compounds


Polar water molecules pull ions away from the
crystal and surround them.
Many ionic compounds are soluble in water.
6-3 Solubility of Ionic Compounds


In solid state, ions
can’t move, can’t
conduct electricity
When ionic
compounds
dissolve in water,
the charged
particles are free
to move – the
solution can
conduct electricity
6-3 Polyatomic Ions


Monatomic ions –
form when a single
atom gains or loses an
electron or electrons
Polyatomic ions – form
when a group of
atoms that are bonded
covalently take on a
charge

ammonium

nitrate

sulfate

carbonate
6-3 Lewis Structures of Polyatomic
Ions/Resonance

Ammonium

Sulfate

Nitrate

Carbonate
6-4 Metallic Bonding


The unique
properties of
metals can be
accounted for by
the metallic bond.
Conduct heat and
electricity,
malleable, ductile,
luster
6-4 Electron Sea Model





Metals have only 1, 2 or 3
valence electrons
Also have vacant p- and
d- orbitals
When metal atoms are
close to each other, these
vacant orbitals overlap
Outer electrons roam
freely throughout network
of overlapping orbitals –
electrons are delocalized
Metallic bonding – results
from attraction between
metal atoms and the
surrounding sea of
electrons
6-4 Conductivity and Luster


When charged
particles are free to
move, an electrical
current can pass
through – metals
conduct electricity
Because metal atoms
have many orbitals
separated by small
energy differences,
metals absorb many
light frequencies –
when energy is
emitted, light is
released – looks shiny
6-4 Malleability and Ductility

When atoms are moved, electrons
flow around them and take new
shape
6-4 Metallic Bond Strength


Varies with nuclear
charge of atoms and
number of electrons in
electron sea (metals
with one valence
electron are softer
than metals with 2
valence electrons)
Heat of vaporization –
heat required to
vaporize a metal is a
measure of the
strength of the bonds
that hold it together
period Heats of Vaporization,
kJ/mol
2nd
Li,
147
Be,
297
3rd
Na,
97
Mg,
128
Al,
294
4th
K,
77
Ca,
155
Sc,
333
5th
Rb,
76
Sr,
137
Y,
365
6th
Cs,
64
Ba,
140
La,
402
Table 6-4 on p. 182
6-5 Molecular Geometry


VSEPR – Valence Shell Electron Pair
Repulsion
VSEPR theory is a model that
accounts for the shapes of simple
molecules
6-5 VSEPR Theory


Repulsion between sets of valence
electrons surrounding an atom
causes these sets to be oriented as
far apart as possible
BeF2
6-5 VSEPR Theory

BF3

CH4
6-5 VSEPR

NH3

H2O
A – central atom

The shapes of
simple
molecules are
determined by
the number of
atoms bonded
to the central
atom and the
number of
unshared pairs
of electrons
around the
central atom.
X – atom bonded to
central atom
E – unshared electron
pair on central atom
6-5 VSEPR

SF2

PCl3
6-5 VSEPR

CHCl3

CO2
6-5 Hybridization



Hybridization is a model that explains how
the orbitals of an atom are rearranged
when the atom forms covalent bonds
Hybridization is the mixing of two or more
atomic orbitals of similar energies on the
same atom to produce new orbitals of
equal energies
Especially useful for explaining bonding in
carbon compounds
6-5 Methane


Methane has tetrahedral geometry
(predicted by VSEPR and known from
experimentation), but valence
electrons of carbon atom are in 2
different kinds of orbitals
How does carbon make four
equivalent covalent bonds in this
compound?
6-5 Methane
___ ___ ___
2p
___ ___ ___ ___
sp3
___
2s
6-5 Methane
6-5 Hybrid Orbitals


Orbitals of equal energy produced by
the combination of two or more
orbitals on the same atom
Number of hybrid orbitals equals
number of atomic orbitals that have
combined
6-5 Hybridization
Atomic
Orbitals
Type of
Hybridization
s, p
sp
Number of
Hybrid
Orbitals
2
s, p, p
sp2
3
s, p, p, p
sp3
4
6-5 Ethane
6-5 Ethene
6-5 Ethyne
6-5 Intermolecular Forces
(van der Waals Forces)



Forces of attraction between
molecules
Vary in strength but generally
weaker than ionic, metallic or
covalent bonds
Boiling point is a good measure of
the strength of intermolecular forces
bonding type
substance
bp (1 atm, °C)
nonpolar-covalent
H2
-253
(molecular)
O2
-183
Cl2
-34
Br2
59
CH4
-164
CCl4
77
C6H6
80
polar-covalent
PH3
-88
(molecular)
NH3
-33
H2S
-61
H2O
100
HF
20
HCl
-85
ICl
97
NaCl
1413
MgF2
2239
Cu
2567
Fe
2750
W
5660
ionic
metallic
Table 6-7, p. 190
6-5 Molecular Polarity and DipoleDipole Forces


Polar molecules (like water) are dipoles.
They have two poles, one positive and one
negative.
Forces of attraction between polar
molecules are called dipole-dipole forces.
6-5 Dipole-Dipole Forces



Short range
Act only between nearby molecules
Polarity of molecules is determined
by types of bonds and arrangement
of bond
6-5 Polarity

Water

ammonia
6-5 Polarity


carbon
tetrachloride
carbon dioxide
6-5 Dipole-Induced Dipole


Electron clouds are mobile
A permanent dipole (like water, ammonia,
hydrogen chloride) can induce a
temporary dipole in a nonpolar molecule
6-5 Hydrogen Bonding



Occurs in compounds in which hydrogen is
attached to oxygen, nitrogen or fluorine
Very strong
Many of water’s special properties can be
accounted for by hydrogen bonding
6-5 London Dispersion Forces




Weak forces
Electrons are in
constant motion
Molecules can have
temporary dipoles
due to this
movement of
electrons
A temporary dipole
can induce another
temporary dipole





Fatty acids can be
saturated of unsaturated.
Saturated fatty acids
have all single bonds.
Unsaturated fatty acids
have some double bonds.
Double bonds cause a
kink in the carbon chain.
Unsaturated fatty acids
don’t pack together as
well, have weaker
dispersion forces
between them, are less
likely to form solid in
arteries.