Unit 2: Chemical Bonding
Chemistry2202
Outline
Bohr diagrams
 Lewis Diagrams
 Types of Bonding
 Ionic bonding
 Covalent bonding (Molecular)
 Metallic bonding
 Network covalent bonding

Types of Bonding (cont’d)
 London Dispersion forces
 Dipole-Dipole forces
 Hydrogen Bonding
 VSEPR Theory (Shapes)
 Physical Properties

Bohr Diagrams (Review)
How do we draw a Bohr Diagram for
- The F atom?
- The F ion?
Draw Bohr diagrams for the atom and
the ion for the following:
Al
S
Cl
Be
Lewis Diagrams
LD provide a method for keeping
track of electrons in atoms, ions, or
molecules
 Also called Electron Dot diagrams
 the nucleus (P& N) and filled energy
levels are represented by the element
symbol

Lewis Diagrams

dots are placed around the element
symbol to represent valence electrons
Lewis Diagrams
eg. Lewis Diagram for F
lone pair
bonding
electron
••
•
•F •
••
lone pair
lone pair
Lewis Diagrams
lone pair – a pair of electrons not
available for bonding
bonding electron – a single electron
that may be shared with another atom
Lewis Diagrams
eg. Lewis Diagram for C
•
•
•C
•
Lewis Diagrams
eg. Lewis Diagram for P
••
•
•P
•
Lewis Diagrams
eg. Lewis Diagram for Na
•
Na
Lewis Diagrams
For each atom draw the Lewis diagram
and state the number of lone pairs and
number of bonding electrons
Li
Be
Al
Si
Mg
N
B
O
Complete bonding worksheet #1
Lewis Diagrams for Compounds
draw the LD for each atom in the
compound
 The atom with the most bonding
electrons is the central atom
 Connect the other atoms using single
bonds (1 pair of shared electrons)
 In some cases there may be double
bonds or triple bonds

Lewis Diagrams for Compounds
eg. Draw the LD for:
PH3
CF4
Cl2O
C2H6
C2H4
C2H2
Lewis Diagrams for Compounds
eg. Draw the LD for:
NH3
SI2
POI
N2
SiCl4 N2H4
CO2
N2H2
CH3OH
H2
O2
HCN
CH2O
Lewis Diagrams for Compounds
A structural formula shows how the
atoms are connected in a molecule.
To draw a structural formula:
 replace the bonded pairs of electrons
with short lines
 omit the lone pairs of electrons
Complete bonding worksheet #2
Why is propane (C3H8) a gas at STP while
kerosene (C10H22) a liquid?
Why is graphite soft enough to write with while
diamond is the hardest substance known even
though both substances are made of pure
carbon?
Why can you tell if it is ‘real gold’ or just ‘fool’s
gold’ (pyrite) by hitting it with a rock?
‘As Slow As Cold Molasses’
‘All Because of Bonding’
‘liquids’ @ -30 ºC
Bonding
Bonding between atoms, ions and
molecules determines the physical and
chemical properties of substances.
Bonding can be divided into two
categories:
- Intramolecular forces
- Intermolecular forces
Bonding
Intramolecular forces are forces of
attraction between atoms or ions.
Intramolecular forces include:
1. ionic bonding
2. covalent bonding
3. metallic bonding
4. network covalent bonding
Bonding
Intermolecular forces are forces of
attraction between molecules.
Intermolecular forces include:
5. London Dispersion Forces
6. Dipole-Dipole forces
7. Hydrogen Bonding
Ionic and Covalent Bonding
ThoughtLab p. 161
Identify #’s 1 - 6
Ionic Bonding



Occurs between cations and anions –
usually metals and non-metals.
An ionic bond is the force of attraction
between positive and negative ions.
Properties:
conduct electricity as liquids and in solution
 hard crystalline solids
 high melting points and boiling points
 brittle

Ionic Bonding
In an ionic crystal
the ions pack tightly
together.
 The repeating 3-D
distribution of
cations and anions
is called an ionic
crystal lattice.

Ionic Bonding
Each anion can be
attracted to six or
more cations at
once.
 The same is true
for the individual
cations.

Ionic Bonding
Covalent Bonding
Occurs between non-metals in molecular
compounds.
 Atoms share bonding electrons to become
more stable (noble gas structure).
 A covalent bond is a simultaneous
attraction by two atoms for a common pair
of valence electrons.

Covalent Bonding
Molecular compounds
have low melting and
boiling points.
 Exist as distinct
molecules.

Covalent Bonding
Molecular
compounds do
not conduct
electric current
in any form
Property
Type of
elements
Force of
Attraction
Ionic
Molecular
Metals and
nonmetals
Non-Metals
Positive ions attract Atoms attract a
negative ions
shared electron
pair
Electron
Electrons move
Electrons are
movement from the metal to
shared
the nonmetal
between atoms
State at room
Always solids
Solids, liquids,
temperature
or gas
Property
Solubility
Ionic
Molecular
Soluble or low Soluble or
solubility
insoluble
Conductivity in None
solid state
None
Conductivity in Conducts
liquid state
None
Conductivity in Conducts
solution
None
Complete bonding worksheet #3
Metallic Bonding (p. 171)
metals tend to lose valence electrons.
 valence electrons are loosely held and
frequently lost from metal atoms.
 This results in metal ions surrounded by
freely moving valence electrons.
 metallic bonding is the force of attraction
between the positive metal ions and the
mobile or delocalised valence electrons

Metallic Bonding
Metallic Bonding

This theory of metallic bonding is called
the ‘Sea of Electrons’ Model or ‘Free
Electron’ Model
Metallic Bonding

This theory accounts for properties of metals
1. electrical conductivity
- electric current is the flow of charge
- metals are the only solids in which charged
particles are free to move
2. solids
- Attractive forces between positive cations and
negative electrons are very strong
Metallic Bonding
3.
-
-
malleability and ductility
metals can be hammered into thin
sheets(malleable) or drawn into thin
wires(ductile).
metallic bonding is non-directional such that
layers of metal atoms slide past each other
under pressure.
Network Covalent Bonding (p. 199)
occurs in 3 compounds (memorize these)
 diamond – Cn
 carborundum – SiC
 quartz – SiO2
 large molecules with covalent bonding in
3-d
 each atom is held in place in 3-d by a
network of other atoms

Network Covalent bonding

Properties:
 the highest melting and boiling points
 the hardest substances
 brittle
 do not conduct electric current in any
form
2. Ionic bonding(metal & nonmetal)
3. Metallic bonding (metals)
4. Molecular (nonmetals)
Weakest
MP & BP decreases
Strongest
1. Network Covalent (Cn ,SiO2 , SiC)
Valence Shell Electron Pair
Repulsion theory (VSEPR)
The shape of molecules is determined by
the arrangement of valence electron pairs
around the atoms in a compound.
 There are 5 shapes that can be
determined by the # of bonds and # of
lone pairs on the central atom.

1. Tetrahedral (4 bonds; 0 lone pairs)
2.
Pyramidal (3 bonds; 1 lone pair)
3. V-shaped (2 bonds; 2 lone pairs)
4. Trigonal Planar (3 bonds; 0 lone pairs)
5. Linear (2 bonds; 0 lone pairs)
For each molecule below draw the
Lewis diagram, structural diagram
shape name.
HOCl
H2Se
H2O2
NBr3
C2F4
C2H6
CHCl3
CH3OH
PBr3
I2
SiH4
HCN
HSiHO
C2H2
Complete molecular models lab
Electronegativity (EN - p. 174)
EN is a measure of the attraction that
an atom has for shared electrons.
 A higher EN means a stronger
attraction or electrostatic pull on
valence electrons
 EN values increase as you move:
- from left to right in a period
- up in a group or family

Increases
The Bonding Continuum
Electronegativity & Covalent Bonds
1. polar covalent bond
a bond between atoms with different EN
- the shared electron pair is attracted
more strongly to the atom with the
higher EN
δ−
δ+
-
H Cl
Electronegativity (p. 174)
polar covalent bond


a covalent bond between atoms with different EN
the shared electron pair is attracted more
strongly to the atom with the higher EN
nonpolar covalent bond

a covalent bond between atoms with the same
EN
bond dipole

An arrow drawn in the direction of the atom with
the greatest EN
Complete:
#’s 7 – 9 on p.178
Weakest
Covalent (nonmetals)
→ London Dispersion
(all molecules)
→ Dipole-Dipole
(polar molecules)
→ H bonding
(H-N, H-O, H-F)
p. 226 #13
Omit parts g), j) – o), q), u), & v)
- Answers on p. 815 for #13
- Incorrect answers
c), d), & s)
Electronegativity and Ionic Bonds

Because the EN of metals is so
low, metals lose electrons to form
cations

Nonmetals gain electrons to form
anions because the EN of
nonmetals is relatively high
Electronegativity and Ionic Bonds

When ions form, the resulting
electrostatic force is an ionic bond
Electronegativity and Covalent Bonds

Atoms in covalent compounds can
either have:


the same EN
eg. Cl2 , PH3, NCl3
different EN
eg. HCl
Electronegativity and Covalent Bonds
Atoms that have the same EN attract
the shared valence electrons to the
same extent.
 Covalent bonds resulting from equal
sharing of the bonding electron pairs
are called Nonpolar Covalent
Bonds

Electronegativity and Covalent Bonds
Atoms that have different EN attract
the shared pair of valence electrons
at different strengths
 The atom with the higher EN exerts
a stronger attraction on the shared
electron pair
eg. H2O

Electronegativity and Covalent Bonds
Since the oxygen atom has a higher EN
the bonding electrons will be pulled
closer to the oxygen atom
 This results in slight positive and
negative charges within the bond.


These charges are referred to as
“partial charges” and are denoted
with the Greek letter delta (δ).
Electronegativity and Covalent Bonds

The region around the oxygen atom
will be slightly negative, and around
the hydrogens will be slightly positive
Electronegativity and Covalent Bonds
The symbol, δ+ represents a partial
positive charge (less than +1) and
δ− represents a partial negative
charge (less than −1).
 Since the bond is polarized into a
positive area and a negative area
the bond has a “bond dipole”.

Electronegativity and Covalent Bonds

p.178
The arrow points to
the atom with the
higher EN.
Electronegativity and Covalent Bonds
 Covalent
bonds resulting from
unequal (electronegativities)
sharing of bonding electron
pairs are called Polar
Covalent Bonds
Electronegativity Homework
#’s 7, 8, & 9 - p. 178
#’s 1, 2, & 3 - p. 180
Bond Energy (pp. 179-180)
1. Describe the forces of attraction and
repulsion present in all bonds.
2. What is bond length?
3. Define bond energy.
4. Which type of bond has the most energy?
5. How can bond energy be used to predict
whether a reaction is endothermic or
exothermic?
Test Outline
Bohr Diagrams (atoms & ions)
 Lewis Diagrams (Electron Dot)
 Ion Formation
 Ionic Bonding, Structures & Properties
 Covalent Bonding, Structures & Properties

Test Outline
Metallic Bonding Theory& Properties
 Network Covalent Bonding & Properties
 Electronegativity
 Bond Dipoles & Polar Molecules
 VSEPR Theory
 LD, DD, & H-bonding
 Predicting properties (bp, mp, etc.)

Molecular Dipoles
The vector sum of all the bond
dipoles in a molecule is a Molecular
Dipole
 A Polar Molecule has a molecular
dipole that points toward the more
electronegative end of the molecule.
eg. H2O

Molecular Dipoles

NonPolar Molecules DO NOT have
molecular dipoles. This occurs when:
- the bond dipoles cancel
- there are no bond dipoles
eg. CO2
PH3
Molecular Dipoles
To determine whether a molecule is
polar:
- draw the Lewis diagram and the
structural diagram
- draw the bond dipoles and determine
whether they cancel

Intermolecular Forces
Strongest bonds; Highest mp and bp
1. Network Covalent (Cn SiO2 SiC)
2. Ionic bonding(metal & nonmetal)
3. Metallic bonding (metals)
4. Molecular (nonmetals)
Weakest bonds; Lowest mp and bp
- Intermolecular forces present
To compare mp and bp in covalent
compounds you must use:
- London Dispersion forces (p. 204)
(all molecules)
- Dipole-Dipole forces (pp. 202, 203)
(polar molecules)
- Hydrogen Bonding (pp. 205, 206)
(H bonded to N, O, or F)
Intermolecular Forces (p. 202)
Intermolecular Forces
Covalent compounds have low mp and
bp because forces between molecules
in covalent compounds are very weak.
 Intermolecular forces were studied
extensively by the Dutch physicist
Johannes van der Waals
 In his honor, two types of intermolecular
force are called Van der Waals forces.

Intermolecular Forces
Intermolecular forces can be used to
account for the physical properties of
covalent compounds.
Intermolecular Forces
Van der
Waals
1. London Dispersion Forces
• LD forces exist in ALL molecular elements
& compounds.
•The positive charges in one molecule
attract the negative charges in a second
molecule.
• The temporary dipoles caused by electron
movement in one molecule attract the
temporary dipoles of another molecule.
1. London Dispersion Forces
The strength of these forces depends on:
a) the number of electrons
more electrons produce stronger LD
forces that result in higher mp and bp
eg. CH4 is a gas at room temperature.
C8H18 is a liquid at room temperature.
C25H52 is a solid at room temperature.
Account for the difference.
1. London Dispersion Forces
Two molecules that have the same number
of electrons are isoelectronic
eg. C2H6 and CH3F
1. London Dispersion Forces
b) shape of the molecule
molecules that “fit together” better will
experience stronger LD forces
eg. Cl2 vaporizes at -35 ºC while C4H10
vaporizes at -1 ºC. Use bonding to
account for the difference.
2. Dipole-dipole Forces
- occur between polar molecules
- the δ+ end of one polar molecule is
attracted to the δ- end of another polar
molecule (& vice-versa)
eg. Which has the higher boiling point
CH3F or C2H6 ?
p. 202
3. Hydrogen Bonds
- a special type of dipole-dipole force
(about 10 times stronger)
- only occurs between molecules that
contain a H atom which is directly
bonded to F, O, or N
ie. the molecule contains at least one
H-F, H-O, or H-N covalent bond.
3. Hydrogen Bonds
-the hydrogen bond occurs between the
H atom of one molecule and the N, O,
or F of a second molecule.
eg. Arrange these from highest to
lowest boiling point
C3H8
C2H5OH
C2H5F
p. 206
NOTE: To compare covalent compounds
you must use:
- London Dispersion forces
(all molecules)
- Dipole-Dipole forces
(polar molecules)
- Hydrogen Bonding
(H bonded to N, O, or F)
Complete bonding worksheet #4
p. 210
Intermolecular Forces
1. Use intermolecular forces to explain the following:
a) Ar boils at -186 °C and F2 boils at -188 °C .
b) Kr boils at -152 °C and HBr boils at -67 °C.
c) Cl2 boils at -35 °C and C2H5Cl boils at 13 °C .
2. Examine the graph on p. 210:
a) Account for the increase in boiling point for the
hydrogen compounds of the Group IV elements.
b) Why is the trend different for the hydrogen
compounds of the Group V, VI, and VII elements?
c) Why are the boiling points of the Group IVA
compounds consistently lower than the others.
3.Which substance in each pair has the higher
boiling point. Justify your answers.
(a)
SiC or KCl
(b)
RbBr or C6H12O6
(c)
C3H8 or C2H5OH
(d)
C4H10 or C2H5Cl
2. Examine the graph on p. 210:
a) Account for the increase in boiling point for
the hydrogen compounds of the Group IV
elements.
b) Why is the trend different for the hydrogen
compounds of the Group V, VI, and VII
elements?
c) Why are the boiling points of the Group IVA
compounds consistently lower than the other
compounds.
Dipole-Dipole Forces

In the liquid
state, polar
molecules
(dipoles) orient
themselves so
that oppositely
charged ends of
the molecules are
near to one
another.
Summary
The types of bonding/forces ranked from
strongest to weakest are:
Strongest
- Network Covalent
- Ionic
- Metallic
Weakest
- Covalent
NOTE: To compare covalent compounds
you must use:
- London Dispersion forces
(all molecules)
- Dipole-Dipole forces
(polar molecules)
- Hydrogen Bonding
(H bonded to N, O, or F)
p. 226 #’s 13 & 14
Dipole-Dipole Forces

The electrostatic attractions
between these oppositely charged
ends of the polar molecules are
called dipole-dipole forces.
Dipole-Dipole Forces
Results of dipole-dipole attractions:
 polar molecules will tend to
attract one another more at room
temperature than similarly sized
non-polar molecules
 energy needed to separate polar
molecules is therefore higher than
for non-polar molecules of similar
molar mass
Dipole-Dipole Forces
Results of dipole-dipole attractions:
 The melting points and boiling
points of substances made of polar
molecules are higher than for
substances made of non-polar
molecules.
Ion-Dipole Forces

An ion-dipole force is the force of
attraction between an ion and a polar
molecule (a dipole).
Ion-Dipole Forces

NaCl dissolves in water because the
attractions between the Na+ and Cl- ions
and the partial charges on the H2O
molecules are strong enough to overcome
the forces that bind the ions together.
Induced Intermolecular Forces
Induced Intermolecular Forces

Induction of electric charge occurs
when a charge on one object
causes a change in the distribution
of charge on a nearby object. (for
example, the balloon)
Induced Intermolecular Forces
There are two types of charge induced dipole forces:
1. An ion-induced dipole force
results when an ion in close
proximity to a non-polar molecule
distorts the electron density of the
non-polar molecule
Induced Intermolecular Forces
The molecule then becomes
momentarily polarized, and the two
species are attracted to each other. (ie.
hemoglobin)
2. In a dipole-induced dipole force
the charge on a polar molecule is
responsible for inducing the charge on
the non-polar molecule.
Dispersion (London) Forces
Bond vibrations, which are part of
the normal condition of a non-polar
molecule, cause momentary,
uneven distribution of charge;
 a non-polar becomes slightly
polar for an instant, and continues
to do so in a random but constant
basis.

Dispersion (London) Forces
At the instant that one non-polar
molecule is in a slightly polar
condition, it is capable of inducing
a dipole in a nearby molecule
 This force of attraction is called a
dispersion force.

Dispersion (London) Forces
Two factors affecting the
magnitude of dispersion forces
are:
1. The number of electrons in the
molecule:
Vibrations within larger molecules
that have more electrons than
smaller molecules can easily
cause an uneven distribution of
charge.

Dispersion (London) Forces
The dispersion forces between these
larger molecules are thus stronger,
which has the effect of raising the
boiling point for larger molecules.
2. The shape of the molecule:
A molecule with a spherical shape has a
smaller surface area than a straight
chain molecule that has the same
number of electrons
Dispersion (London) Forces
Therefore, the substance with
molecules that have a more
spherical shape will have weaker
dispersion forces and a lower
boiling point.
 London dispersion forces are
responsible for the formation and
stabilization of the biological
membranes surrounding every
living cell.
Hydrogen Bonding
In order to form a hydrogen bond, a
hydrogen atom must be bonded to
a highly electronegative atom such
as oxygen, nitrogen, or fluorine.
Hydrogen Bonding
These bonds are very polar, and
since hydrogen has no other
electrons, the positive proton, H+,
is exposed and can become
strongly attracted to the negative
end of another dipole nearby
Hydrogen Bonding
A hydrogen bond is an
electrostatic attraction between the
nucleus of a hydrogen atom,
bonded to fluorine, oxygen, or
nitrogen and the negative end of a
dipole nearby.
δ+
δ−
δ+
δ+
δ−
…
Hydrogen Bonding
In biological systems, these polar
bonds are often parts of much
larger molecules (ie. N H bonds
and H O bonds found in biological
molecules)
Hydrogen Bonding in Water
Hydrogen bonds between the
hydrogen atoms in one water
molecule and the oxygen atom in
another account for many unique
properties of water.
δ+
δ−
δ+
δ+
δ−
…
Hydrogen Bonding in Water


In liquid water, each
water molecule is
hydrogen bonded to
at least four other
water molecules.
The large number
of bonds between
water molecules
makes the net
attractive force quite
strong
Hydrogen Bonding in Water
the strong attractive forces are
responsible for the relatively high
boiling point of water.
 The water molecules are farther
apart in ice then they are in liquid
water making ice less dense than
liquid water.

Hydrogen Bonding in Water

Hydrogen bonds
force water
molecules into the
special hexagonal,
crystalline structure
of ice when the
temperature is
below 4 degrees
celcius.